AP Chemistry Review: Unit 1 (Atomic Structure and Properties)

Summary

Professor Dave provides a detailed review of Unit 1 of the AP Chemistry curriculum, focusing on atomic structure and properties. This segment covers key concepts such as electron configurations, periodic trends, isotopes, ions, and spectroscopy. Starting from the basic structure of atoms and progressing through complex ideas like quantum numbers and periodic trends, this overview aims to simplify the understanding of chemistry's fundamental building blocks, ensuring students are well-prepared for the AP exam.

Highlights

• Atomic structure is the foundation of chemistry, including moles, isotopes, and ions 🌟.
• Electron configurations are determined by energy levels and filling orders 🎯.
• Periodic trends like atomic radius and ionization energy are crucial for understanding chemical behavior 📊.
• Spectroscopy techniques provide insights into isotopic and elemental properties 🔍.
• Quantum numbers provide a detailed understanding of electron distributions within atoms ⚛️.

Key Takeaways

• Understanding atomic structure and properties is crucial for AP Chemistry success 📘.
• Valence electrons determine reactivity and are key to periodic table trends 🔍.
• Key periodic trends include atomic radius, ionization energy, electron affinity, and electronegativity 📈.
• The Aufbau principle and Hund's rule guide electron configuration filling 📜.
• Mass spectroscopy helps define isotopic composition and abundance in elements 🔬.

Overview

In this comprehensive review, Professor Dave demystifies the basics of atomic structure and properties, making it user-friendly for AP Chemistry students. From the fundamental concepts like moles and molar mass to more intricate ideas like spectroscopy and periodic trends, this video serves as an ideal study companion.

As students navigate through the video, they will gain insight into how atoms form molecules, how the periodic table is structured and why it’s organized in its particular way. Essential concepts like atomic mass, electron configurations, and chemical bonds are broken down into digestible bits, which students can easily understand and recall.

The video closes with a focus on periodic trends, a crucial concept in chemistry that describes the predictable patterns in element behavior. By the end of this session, students are well-equipped with the crucial knowledge of Unit 1, ready to take on more complex chemical topics in their AP journey.

Chapters

• 00:00 - 00:30: Introduction: Preparing for AP Chemistry Exam The introduction sets the stage for preparing for the AP Chemistry exam, emphasizing the importance of understanding various terminologies and concepts. Dave Farina indicates the approach to cover all necessary terms and ideas in a condensed format through comprehensive videos, each focusing on one of the nine units covered in the exam. The introduction then transitions into discussing Unit 1, which deals with atomic structure and properties.
• 00:30 - 01:00: Topics Overview of Unit 1 The chapter covers fundamental chemistry concepts, including moles and molar mass, mass spectroscopy, pure substances and mixtures, electron configurations, photoelectron spectroscopy, periodic trends, valence electrons, and ionic compounds. The focus will be on complex topics as they arise, but reviewing the fundamentals is encouraged.
• 01:00 - 01:30: Basic Concepts: Atoms and Molecules This chapter introduces the basics of atoms and molecules, starting with a definition of an atom as the smallest unit of an element. The text uses copper as an example, explaining how repeatedly dividing a copper rod eventually leads to the atomic level, where the material still retains its identity as copper. The uniqueness of elements is attributed to their distinct atomic structures, particularly the number of protons present. Furthermore, it touches upon how atoms combine through chemical bonds to form molecules.
• 02:00 - 03:30: Pure Substances and Mixtures This chapter explains the difference between elements, compounds, and mixtures. Elements can be single atoms or molecules made of the same type of atoms, like diatomic oxygen (O2) and nitrogen (N2). Compounds are formed when atoms of different elements bond together, such as carbon dioxide (CO2) made from one carbon atom and two oxygen atoms, and water (H2O) made from one oxygen atom and two hydrogen atoms.
• 04:00 - 06:00: Atomic Structure The chapter titled 'Atomic Structure' discusses the concept of pure substances, which can be either elements or compounds. It explains that a pure substance is composed of only one type of particle. For instance, sodium metal is an element and thus a pure substance because it consists solely of sodium atoms. Similarly, water, despite being a compound composed of hydrogen and oxygen elements, is also a pure substance because it consists only of water molecules. The chapter underscores that although a water molecule is composed of more than one type of atom, its identity and properties as a water molecule are consistent.
• 11:00 - 13:00: Moles, Atomic Mass, and Molar Mass The chapter titled 'Moles, Atomic Mass, and Molar Mass' explains the concept of mixtures and compounds in chemistry. It highlights the differences between them, noting that mixtures consist of different substances with separate identities and properties, whereas compounds are composed of atoms from different elements that form a single molecule with chemical bonds. The chapter details how these compositions determine the macroscopic properties of materials, using water as an example of a compound.
• 14:30 - 19:00: Compounds, Percent Composition, and Formulas The chapter discusses the concepts of compounds, percent composition, and chemical formulas. It explains the differences between mixtures and compounds, with a focus on homogeneous and heterogeneous mixtures. A homogeneous mixture, such as sucrose dissolved in water, is one in which the substances involved are evenly distributed within the container. This means that any portion of the mixture would look the same if examined closely, with molecules of each substance dispersed uniformly.
• 19:30 - 25:30: Quantum Numbers and Atomic Orbitals The chapter 'Quantum Numbers and Atomic Orbitals' seems to be about fundamental chemistry concepts, likely elaborating on atomic structure and the distribution of electrons in atoms using quantum mechanics. The excerpt provided contrasts homogeneous and heterogeneous mixtures, using oil and water as examples of a heterogeneous mixture. It suggests that while homogeneous mixtures have a uniform composition, heterogeneous mixtures like oil and water maintain distinct layers that do not mix even at the interface.
• 26:00 - 28:00: Electron Configurations This chapter provides a summary of the foundational concepts in chemistry, focusing on the classification of matter. It differentiates between pure substances, which include elements and compounds, and mixtures, which can be further classified as homogeneous or heterogeneous based on their uniformity. The chapter sets the stage for discussing atomic structures, emphasizing the arrangement and roles of protons, neutrons, and electrons in forming atoms.
• 28:30 - 36:30: Periodic Trends This chapter explores the internal structure of an atom, focusing on the relative positions and masses of protons, neutrons, and electrons. It highlights that protons and neutrons, which are located in the nucleus, have approximately the same mass, while electrons, located far from the nucleus, have significantly less mass.
• 39:00 - 39:30: Conclusion and Transition to Unit 2 In this chapter, the fundamental charges of protons and electrons are discussed. Protons carry a positive charge, and electrons carry an equal but negative charge, with each unit of charge being 1.602 x 10⁻¹⁹ coulombs. Neutrons are neutral as they have no charge. The identity of an element is defined by the number of protons in its nucleus, illustrated through the example of carbon, which always has six protons. The chapter transitions towards the next unit, suggesting further exploration of atomic structure and properties.

AP Chemistry Review: Unit 1 (Atomic Structure and Properties) Transcription

• 00:00 - 00:30 Hey everyone, Dave Farina here. To get ready  for the AP chemistry exam, we’re going to have   to understand a lot of different terminology  and concepts, and be able to answer fairly   complicated questions about them. Because  there are so many terms and concepts to know,   let’s go through them all together in as condensed  a format as possible, with one comprehensive video   for each of the nine units the exam will cover. Unit 1 is on atomic structure and properties, so
• 00:30 - 01:00 these will be all the very basics. We will discuss  moles and molar mass, mass spectroscopy, pure   substances and mixtures, electron configurations,  photoelectron spectroscopy, periodic trends,   valence electrons, and ionic compounds. We will be focusing on the more complicated   topics as they arise but it will be a good  idea to quickly review the fundamentals,   so let’s start with an abbreviated  version of some basic concepts.
• 01:00 - 01:30 An atom is the smallest unit of an element, which  are these things we see on the periodic table,   such as copper. If we cut a copper rod in half,  we get two smaller copper rods. We can continue   doing this many times until we finally reach the  tiniest thing that could be considered copper,   and that’s a copper atom. Every element is made  of a different type of atom which means it has   a different number of protons inside, and  if we were to break apart this copper atom,   we wouldn’t have copper anymore. Atoms join together via chemical
• 01:30 - 02:00 bonds to form molecules. Sometimes molecules can  be considered elements, if they are comprised of   only one type of atom, like diatomic oxygen and  nitrogen. But when atoms of more than one element   come together to form a molecule, we would call  this a compound. One carbon atom and two oxygen   atoms makes carbon dioxide. One oxygen atom  and two hydrogen atoms makes a water molecule.
• 02:00 - 02:30 A pure substance can be an element or compound.  If we have a block of sodium metal, which is an   element, that’s a pure substance, because it  is made of just one thing, sodium atoms. If we   have a sample of water, which is a compound,  that’s a pure substance, because it is also   made of just one thing, water molecules. Each  water molecule has more than one element in it,   hydrogen and oxygen, but it is still a water  molecule, and the properties of water molecules
• 02:30 - 03:00 are what determine the macroscopic properties  of water. When we put more than one element or   compound together, we get a mixture. Notice that  we have two different substances in this mixture,   with totally separate identities and properties.  So a compound is a combination of atoms of   different elements to form a single molecule,  where all the atoms are chemically bound, like   with water. But with a mixture, we have substances  that are totally different elements or compounds,
• 03:00 - 03:30 simply mixed together in the same space. Mixtures can be either homogeneous or   heterogeneous. A homogeneous mixture involves  substances that, when mixed together, are   distributed evenly in the container. For example,  sucrose dissolved in water. When you put sugar   in water and stir, it will disperse evenly within  the water, such that if you were to zoom in on any   portion of this mixture, it would look exactly the  same, with sugar and water molecules moving around
• 03:30 - 04:00 one another. By contrast, a heterogeneous mixture  will not have its components spread out evenly,   so every section will not look the same. An  example of this is a mixture of oil and water.   We can clearly see that they will form distinct  layers, such that if we zoom in on one part,   we see only oil. If we zoom in on another  part, we see only water. And if we zoom   in at the interface, we see oil and water  molecules up against each other but not mixing.
• 04:00 - 04:30 So to summarize, we have pure substances and  mixtures. A pure substance can be either an   element or a compound. A mixture is  made when more than one type of pure   substance is mixed together, and this can  be either homogeneous or heterogeneous.  Next we are going to want to review atomic  structure, that is the arrangement of protons,   neutrons, and electrons to form atoms,  and the parameters they determine.
• 04:30 - 05:00 Looking inside an atom, the protons and neutrons  sit in the nucleus, with the electrons existing   far away from the nucleus. The proton and neutron  have roughly the same mass, about 1.67 x 10-24   grams, which is a trillionth of a trillionth  of a gram, while the electron has nearly 2,000   times less mass than that, at around 9.1 x  10-28 grams. And in terms of electric charge,
• 05:00 - 05:30 the proton and electron each hold the fundamental  unit of charge, which is 1.602 x 10-19 coulombs,   though for the proton that will be positive,  and for the electron it will be negative.   The neutron is neutral meaning it has no charge. Each element is defined by the number of protons   in its nucleus. For example, a carbon atom has  six protons in its nucleus. Every carbon atom   in the universe has six protons in its nucleus,  and every atom that does not have six protons is
• 05:30 - 06:00 not carbon. This is fundamentally how we define  carbon. So carbon has an atomic number of six,   because the atomic number of an atom is equal  to the number of protons in its nucleus. Every   element has its own unique atomic number, and  therefore characteristic number of protons.   The mass number of an atom is equal to the  sum of the numbers of protons and neutrons,   since each nucleon has a mass of approximately  one atomic mass unit. However, although every
• 06:00 - 06:30 element has a particular atomic number,  elements do not have a specific mass number,   because unlike the number of protons, the number  of neutrons in the nucleus can vary for any given   element. For example, carbon typically has six  neutrons, which combined with the six protons   will produce a mass number of 12. But it can also  have seven neutrons, for a mass number of 13,
• 06:30 - 07:00 or eight neutrons, for a mass number of 14. Atoms  of a given element with different numbers of   neutrons are called isotopes of a given element,  so carbon atoms will always have six protons,   but different isotopes of carbon can have six,  seven, or eight neutrons, which correspond to   different mass numbers. The atomic mass of an  element, which is the other number on each block   of the periodic table, is the average of all the  naturally-occurring isotopes of that element with
• 07:00 - 07:30 respect to their relative abundance. So it’s  essentially the average mass of an atom of that   element. Most carbon is carbon-12, with a little  bit of carbon-13, and trace amounts of carbon-14,   therefore the average mass of all carbon  atoms will be just a tiny bit above 12,   as this equation determines. We just multiply  each mass number by a fraction of one representing
• 07:30 - 08:00 that isotope’s relative abundance, and add them  all together. This can be done for any element.  Finally, since protons are positively charged  and electrons are negatively charged, a neutral   atom will have the same number of protons and  electrons. If an atom gains electrons, it will   become negatively charged, because the number of  negatively charged particles will outnumber the   positively charged ones. And if an atom loses  electrons, it will become positively charged,
• 08:00 - 08:30 because the number of positively charged  particles will outnumber the negatively   charged ones. When an atom is not electrically  neutral, it is called an ion, either a cation   or anion if positive or negative, respectively. With both isotopes and ions understood, we can   understand how the masses of existing isotopes  and their relative abundances can be determined   for any element using a technique called mass  spectroscopy. This instrument takes a sample,
• 08:30 - 09:00 vaporizes it, ionizes it, and sends it through a  tube where it is subjected to an external magnetic   field. The particles then have their paths  deflected by a degree that depends on their mass   to charge ratio, and this information is received  when they collide with a detector. A mass spectrum   displays this mass to charge ratio, which for  ions with a singular charge is essentially just   atomic mass, against their relative abundance. So  with a spectrum like this we can clearly see the
• 09:00 - 09:30 naturally-occurring isotopes for this particular  element with their respective mass numbers on the   horizontal axis, and their relative abundances on  the vertical axis. This is how this information   was initially determined for every element. To quickly summarize, the atomic number of an atom   is equal to the number of protons in the nucleus.  The mass number of an atom is equal to the number
• 09:30 - 10:00 of protons plus the number of neutrons in the  nucleus. This means that we can calculate the   number of neutrons in any atom by finding the mass  number minus the atomic number. And the electrical   charge on an atom is the number of protons minus  the number of electrons. We can report these   values using something called a nuclide symbol.  These consist of the chemical symbol for the   element, with the atomic number in subscript to  the left, the mass number in superscript to the
• 10:00 - 10:30 left, and the charge in superscript to the right.  We will have to interpret these nuclide symbols,   so just remember the definitions of these terms  and where they go. For example, look at this   nuclide symbol for magnesium. We see 12 down  here, which means 12 protons, which is actually   slightly redundant, because every magnesium atom  in the universe must have 12 protons to qualify   as a magnesium atom. Then we see 24 up here,  which means there must be 12 neutrons, since
• 10:30 - 11:00 12 plus 12 is 24. And then this plus two charge  means the atom has lost two electrons compared   to the neutral atom, which leaves 10 electrons. With the submicroscopic structure of the atom   understood, we need to be able to talk about  matter on the macroscopic level, as well. This   is hard, since molecules are way too small to see,  and macroscopic amounts of a substance contain   an unbelievable number of molecules. That’s  why we came up with the concept of the mole.
• 11:00 - 11:30 A mole is just a word that describes a number,  like the way a dozen means 12. But it is a   very large number that allows us to convert  between atomic mass and grams. For example,   carbon weighs on average about 12 atomic mass  units. According to the definition of a mole,   a mole of carbon atoms will therefore weigh 12  grams. This demonstrates how the mole is our
• 11:30 - 12:00 way of converting between atomic mass units  and grams, so that we can discuss molecules   in terms of numbers, but have that number be so  large that it represents a quantity that we can   see with our eyes and do chemistry with. In other  words, this way we can weigh out matter in grams   and do chemistry with it, but still be talking  about numbers of molecules and thus respect the   ratios in which these molecules react. The number  of items in a mole is called Avogadro’s number,
• 12:00 - 12:30 which is equal to 6.022 x 1023, which is nearly a  trillion trillion. So that’s precisely the number   of carbon atoms in 12 grams of carbon, because  one carbon atom weighs 12 atomic mass units.  The mass of one mole of a substance is  called the molar mass. Elements will   have a molar mass equal to their atomic mass  but in grams per mole instead of atomic mass   units. Compounds will also have a molar mass,  and it will be equal to their molecular mass,
• 12:30 - 13:00 so to find the molar mass of a compound we simply  add up the atomic masses of all the atoms in the   molecule, and then we express that number in  grams per mole instead of atomic mass units.  It should be very easy to convert between grams  and moles for any compound. Let’s say we want   to know the number of moles in 28.35 grams of  glycine. We can simply find the molecular mass
• 13:00 - 13:30 by adding up the atomic masses of all the atoms  in the molecule. Using the molecular formula of   C2H5O2N, we get two times twelve for the two  carbons, five times one for the hydrogens,   two times sixteen for the two oxygens, and  fourteen for the lone nitrogen. Adding those   values up will give us a mass of 75, which  if expressed in atomic mass units will be   the molecular mass, but if expressed in grams  per mole it will represent the molar mass,
• 13:30 - 14:00 or the mass contained in one mole of glycine  molecules. Then we convert our mass into moles.   We will multiply our gram value by this conversion  factor, putting moles on the top and grams on   the bottom so that grams cancel, and doing the  arithmetic we will get 0.38 moles of glycine.  We can go the other way as well, from moles to  grams. Let’s say we are looking at vitamin C,
• 14:00 - 14:30 which has the molecular formula C6H8O6. Say  we have 1.42 x 10-4 moles, but we need this   in grams. Once again, adding up the atomic  masses of all the atoms in the molecule,   we can get a molar mass of 176 grams per mole. If  we multiply our value in moles by this conversion   factor, we can see that moles cancel, and we  will get an answer of 0.025 grams of vitamin C.
• 14:30 - 15:00 With moles understood, we can start to better  understand compounds and their compositions,   as well as how these compositions  can be determined. This will involve   discussing empirical and molecular formulas. The percent composition of a compound is the   percent of the molecular mass that is represented  by each element in a compound. This is easy to   calculate if we know the molecular formula of a  compound, because we can then know the molecular
• 15:00 - 15:30 mass. If we know the molecular mass, we can just  find the fraction of the molecular mass that is   contributed by each element. Let’s say we want  to know the percent composition of ammonia. We   know that ammonia has a molecular mass of 17  atomic mass units, because the nitrogen atom   has a mass of 14 atomic mass units, and each  hydrogen atom has a mass of 1 atomic mass unit,   for a total of 17. We can simply calculate the  mass of each element present in the compound
• 15:30 - 16:00 over the total mass of the compound to get  the percent composition of the compound. If   one nitrogen atom is 14, then 14 over 17 will give  us 0.82, which times 100 gives us 82%. This means   that the nitrogen atom in ammonia represents 82%  of the mass of the molecule. The three hydrogen   atoms have a total mass of 3, and 3 over 17 gives  us 0.18. Multiplying by 100, that gives us 18%,
• 16:00 - 16:30 so hydrogen represents 18% of the mass of the  molecule. And 82% plus 18% does add up to 100%,   so these calculations do make sense. This line of thinking is actually a   great way to determine the molecular formula of an  unknown compound. We can do this by first figuring   out the empirical formula of a compound, which  is the lowest whole number ratio of the number
• 16:30 - 17:00 of atoms of different elements in a compound.  Let’s say we combusted an unknown hydrocarbon,   which is a compound consisting of only carbon  and hydrogen, and collected the resulting carbon   dioxide and water. After performing some basic  calculations, we determine that there was 1.71   grams of carbon and 0.287 grams of hydrogen in the  initial sample. Since these values are in grams,   they do not tell us anything about the empirical  formula, because every element has a different
• 17:00 - 17:30 mass. Instead, we must convert these to moles  to make sense of a numerical ratio. We can use   the molar masses of each element to convert to  moles. 1.71 grams of carbon times 1 mole over   12.01 grams gives us 0.142 moles of carbon atoms  in the original sample. Doing the same thing for   hydrogen, 0.287 grams of hydrogen times 1 mole  over 1.008 grams gives us 0.284 moles of hydrogen
• 17:30 - 18:00 atoms in the original sample. Let’s divide both  of these numbers by the smaller number so that   we can try to get a whole number ratio. 0.142 over  0.142 gives us 1, and 0.284 over 0.142 gives us 2,   or a 1 to 2 ratio. So we can see from these  calculations that there must have been twice   as many hydrogen atoms as carbon atoms  in the sample. This makes the empirical
• 18:00 - 18:30 formula for the unknown substance, CH2. We must  realize that this is not the molecular formula,   which tells us the actual number of atoms of each  element in the compound. The compound could have   many more than one carbon atom, but however many  carbon atoms are in the compound, there must be   twice as many hydrogen atoms. At any rate, we  can perform a calculation like this for any   compound containing any combination of elements,  we just use the molar mass of each element and
• 18:30 - 19:00 convert the mass into moles to find the molar  ratios, and therefore the empirical formula.  We can also get the molecular formula if we have  the molecular mass of the compound, which we can   get through mass spectrometry. In such a case,  we would just find out how many multiples of   the formula unit are required to get a total  mass equivalent to the molecular mass. Given   the previous example with an empirical formula of  CH2, let’s say that we knew the molecular mass was
• 19:00 - 19:30 42. The mass of CH2 is 14, 12 from carbon and two  from hydrogen, and 42 divided by 14 is three, so   we just multiply the formula unit by three to get  C3H6, and that must be the molecular formula, as   it obeys the ratio of the empirical formula, and  has a mass that is equal to the molecular mass.
• 19:30 - 20:00 With empirical and molecular formulas  understood, we have to dive back into   the atom and learn more about electrons.  Precisely how are these distributed within   an atom? This will be important to understand  in order to discuss chemical reactions.  Again, atoms contain both protons and electrons,  and these have opposite charge, which means they   are attracted to one another. This attraction  is described mathematically by Coulomb’s law,   which says that the force between two charged  particles is proportional to the product of their
• 20:00 - 20:30 charges divided by the square of the distance  between them. If opposite charges this will be   an attraction, if the same charge it will be  a repulsion. Greater magnitude of charge means   greater force, and closer together means greater  force, while farther away means lesser force.  Electrons themselves reside in things called  atomic orbitals, or three-dimensional regions   of probability surrounding the nucleus where an  electron can be found, and there are different
• 20:30 - 21:00 kinds of quantum numbers which will describe  these orbitals. The first quantum number is   the principal quantum number n. This refers to  the energy level or the shell that the electron   resides in. A higher n value means a higher  energy and further away from the nucleus. The   next number will be the angular momentum quantum  number, L. This can have any value from 0 to n-1,   meaning if n is 1, L is 0. If n is 2, L can be  0 or 1, and so forth. L will define the type
• 21:00 - 21:30 of orbital the electron is in. An L value of 0  corresponds to s orbitals. Those are spherical,   and they increase in radius as n increases. If  L is 1, we are discussing p orbitals, which are   lobes that extend on each of the X, Y, and Z  axes. If L is 2, we are looking at d orbitals,
• 21:30 - 22:00 which look a bit stranger. S, p and d are the  most important ones for our purposes. Next,   we have the magnetic quantum number, m sub L.  This can be anywhere from –L to L, so if L is 2,   and we are discussing d orbitals, L can be -2,  -1, 0, 1, or 2. This is why there are five d   orbitals per energy level, because there are five  possibilities for m sub L and each one corresponds
• 22:00 - 22:30 to an individual orbital. For precisely the  same reason, there are three p orbitals per   energy level, with m sub L values of -1, 0, and 1,  and there is only one s orbital per energy level,   with an m sub L value of zero. Lastly, there is  the spin quantum number, m sub s. This will be   positive one half or negative one half, and  since a maximum of two electrons can fit in   any atomic orbital, each pair will receive  opposite spin values, which we can call spin
• 22:30 - 23:00 up or spin down. The key thing to understand is  that the n value describes a shell of electrons,   and the L value describes a subshell. So there  is an n = 3 shell, and within that there is a 3s   subshell, and a 3p subshell, and a 3d subshell.  Then m sub L describes an individual orbital   within a subshell, and m sub s differentiates  between the two electrons within an orbital.
• 23:00 - 23:30 Now we need to understand how electrons  fill up these orbitals. As n increases,   the energy of the orbital increases, as we are  moving farther away from the nucleus, so Coulomb’s   law says the attraction to the nucleus will  decrease. We should also know that within a shell   the energy increases from s to p to d orbitals.  So 1s is the lowest energy orbital, then 2s,
• 23:30 - 24:00 2p, 3s, 3p, and so forth. But this pattern isn’t  followed precisely when we get to larger atoms,   the first deviation being that the 3d orbitals  are higher in energy than the 4s. Looking at   this diagram, we can see the precise order of the  orbitals in terms of increasing energy. Since a   system will always want to be at the lowest energy  possible, this is the order, from bottom to top,
• 24:00 - 24:30 that an atom will arrange its electrons. This  order in which the orbitals are filled is   called the Aufbau principle. Additionally there is  Hund’s rule, which says that when looking at a set   of degenerate orbitals, which means orbitals of  the same energy, as a set of p orbitals or a set   of d orbitals will always be, we must place one  electron in each orbital first before doubling   them up. So for these p orbitals, each one gets  a spin up electron first, and then we start
• 24:30 - 25:00 generating pairs by placing spin down electrons. Each electron within an atom must be assigned an   orbital, and the specific arrangement of  electrons amongst the orbitals within an   atom is called the electron configuration  of the atom. Many properties of an element   will depend on its electron configuration, so  let’s make sure we understand these as well.  The convention for reporting an electron  configuration is to list all the types of orbitals   that are occupied along with a number to indicate  the occupancy of those particular orbitals. Each
• 25:00 - 25:30 item in an electron configuration should have the  n value, followed by the letter that corresponds   to the type of orbital, and a superscript that  describes the number of electrons contained in   that subshell. So this would be read 2p4, which  refers to a total of 4 electrons that exist in the   2p orbitals. Let’s make sure we understand that  neither of these two numbers is telling us how
• 25:30 - 26:00 many orbitals are being described, as that number  is implied, since each energy level contains 1   s orbital, 3 p orbitals, and 5 d orbitals. So when assigning an electron configuration   we are starting with the lowest energy orbital,  the 1s, and building up to the higher energy   orbitals according to the Aufbau principle and  Hund’s rule, until all the electrons are assigned   to an orbital. A convenient way to follow the  Aufbau principle is to simply know what sections
• 26:00 - 26:30 on the periodic table correspond to which  subshells. Looking at the periodic table now,   we must understand that each period, or  row on the table, represents a shell,   or a particular n value. The first row  is n = 1, then n = 2, and so forth. Then,   we must know that this section containing groups  1 and 2 is called the s block. This section is the   p block. The transition metals are the d block,  and the lanthanides and actinides are the f block,
• 26:30 - 27:00 though we won’t be too concerned with f orbitals  here. The only trick is that the d block is always   one behind the period number in terms of principal  quantum number. For example we can see that in the   4th period it’s actually the 3d orbitals that  follow the 4s. If we can internalize this way   of looking at the periodic table, then the Aufbau  principle reveals itself as we simply read left   to right and up to down on the table. Starting  at the top left corner, the order would be 1s,
• 27:00 - 27:30 2s, 2p, 3s, 3p, 4s, 3d, 4p, and so forth, which  is precisely the order dictated by the Aufbau   principle. This will make it much easier for us  to assign electron configurations to any atom.  We should now be able to assign the electron  configuration of any element. This is easy to   do if we simply look at where an element sits on  the periodic table, and list off the subshells
• 27:30 - 28:00 that the element will utilize by going left to  right and up to down on the table until we get   to the element in question. Take chlorine for  example. In order to get to chlorine, let’s   start up at the top left corner, and read off the  subshells. 1s will be full, so 1s2. Same with 2s,   so 2s2 and 2p6 from period 2. Then 3s2, and  when we get to the 3p orbitals, we count 1,
• 28:00 - 28:30 2, 3, 4, 5 to get to chlorine. 3p5. That gives  us 1s22s22p63s23p5 as the electron configuration   for chlorine. This makes perfect sense, as  neutral chlorine has 17 electrons to place,   and this is the lowest energy configuration  for the distribution of these 17 electrons.  With electron configuration understood, we  can begin to truly understand the structure
• 28:30 - 29:00 of the periodic table, and why elements are  arranged in the groups as they are. Let’s   take a closer look at the periodic table  and see what else we can learn from it.  Looking at the table now, we can see rows called  periods and columns called groups, and elements   will be in the same group because they have  similar electron configurations. Specifically,   they have the same number of valence electrons,  which are the electrons in the outermost shell.
• 29:00 - 29:30 In group 1, all the configurations end in s1. In  group 2, they end in s2. You can see that in every   group, the electron configurations end the same  way as the other elements in that group. So as we   move forward and learn about the periodic table,  it is the number of valence electrons that will   determine the reactivity and properties of any  particular element, as the valence electrons are   the ones that are available to do chemistry. Those  electrons that are not in the outermost shell,
• 29:30 - 30:00 and are therefore not valence electrons, are  called core electrons. These are the ones in   the inner shells, which do not participate  in chemistry. So most elements have many core   electrons, and just a few valence electrons,  particularly as we get lower on the table.  We must comprehend a set of periodic trends,  meaning properties that atoms possess which change   in a predictably periodic way as we move in some  direction along the table. The first property we   will look at is the atomic radius. It’s difficult  to measure the radius of a lone atom, so the
• 30:00 - 30:30 convention is to tabulate lists of covalent radii,  which are defined as one half the distance between   the nuclei of two identical atoms that are bonded  to each other. This should be roughly the same as   the radius of the atom outside of the context  of a chemical bond as well. When we examine   various covalent radii, we notice that the radius  will increase as we go down the periodic table.   This is because when we go down a row on the  table, we are increasing the n value by 1,
• 30:30 - 31:00 thus adding a shell and placing the valence  electrons farther away from the nucleus. This   makes the atom larger, as well as its covalent  radius. Then moving horizontally, as we move to   the right along a period, the covalent radius will  decrease slightly. This may seem counterintuitive,   as the addition of electrons doesn’t seem like  it should result in a smaller radius, but we   must realize that as we move to the right we are  also adding protons, given that the atomic number
• 31:00 - 31:30 is increasing, and the more protons there are  in the nucleus, the greater the electromagnetic   attraction that will pull the electrons in the  existing shells a bit closer to the nucleus. There   are a few deviations to this trend, particularly  if we were to examine the transition elements of   a particular period, but in general, atomic  radius will decrease as we move through a   period. So radius increases going down and left,  and decreases going up and right on the table.
• 31:30 - 32:00 So that covers covalent radii, which essentially  refer to the size of an atom. But what happens   to this radius when an element loses or gains  electrons to become an ion? Any change in the   number of electrons should affect the radius  in some way. As it happens, any time an atom   loses an electron, the remaining valence electrons  will still feel the same effective nuclear charge,   but distributed amongst fewer electrons,  so it will cause the radius to contract.   This means that any cation has a smaller ionic  radius than the covalent radius of the neutral
• 32:00 - 32:30 atom. This difference can be dramatic if all the  valence electrons are lost, since this will result   in the removal of an entire shell, dropping  down to the shell below. But if an atom gains   electrons, the effective nuclear charge will be  distributed amongst more electrons than before,   and there will be additional electron  repulsion amongst the valence electrons,   which results in an expansion of the radius. We  can see a concrete example with aluminum. If all
• 32:30 - 33:00 three valence electrons are lost, it will lose  its entire valence shell, and therefore will have   a dramatically reduced radius. Sulfur on the  other hand, when it becomes the sulfide ion,   the two additional electrons will cause the  radius to increase quite a bit, since there   are no additional protons to pull the electrons,  just more electrons that will push the radius out.  We might sometimes compare isoelectronic  species. These are atoms and ions that have
• 33:00 - 33:30 the same electron configuration. For example,  let’s look at the different species that can   exhibit the electron configuration 1s22s22p6.  We can see that beyond just neon, a number of   different ions can have this configuration if the  corresponding neutral atoms gain or lose a certain   number of electrons. When comparing isoelectronic  species, the radius will decrease as the atomic   number increases. This is because they will all  experience the same amount of electron repulsion,
• 33:30 - 34:00 since they have the same number of electrons,  but as we add more protons to the nucleus,   this pulls the electrons closer to the nucleus. Next let’s look at ionization energy. This is   defined as the energy required to remove the  outermost electron from an atom in the gas phase   and in its ground state configuration. The higher  the ionization energy, the more difficult it is   to remove the electron, which tells us something  about the atomic radius of the atom as well as
• 34:00 - 34:30 the effective nuclear charge felt by the electron.  Each element will have a first ionization energy,   which is the energy required to generate the  1+ cation, and they will also have successive   ionization energies, like the second ionization  energy, to go from 1+ to 2+, and so forth. Each   ionization energy will be larger than the last,  because it will get harder and harder to remove   electrons the more positive the ion becomes, as  each ionization is a further destabilization.
• 34:30 - 35:00 The electron that is removed will always be the  outermost electron. As the atom gets larger,   the outermost electron gets farther away  from the nucleus, and therefore becomes   easier to remove. Every time we add a shell,  we are moving further away from the nucleus,   so ionization energy decreases as we move down the  periodic table. Since atomic radius also decreases   going to the right within a period, we can expect  the ionization energy to increase at the same
• 35:00 - 35:30 time. As we go, we are adding protons, contracting  the radius, and holding electrons more tightly,   so they are harder to remove. That means in  general, while atomic radius increases down   and left, ionization energy will increase up and  right, precisely the opposite of the radius trend.   That means helium is the most difficult element  to ionize, with a single shell that is totally   full and close to the nucleus, while francium is  the easiest, with a lone electron in an outermost
• 35:30 - 36:00 shell that is very far from the nucleus. The energies of the electrons in an atom   can be determined by photoelectron spectroscopy.  With this type of spectrum, the energy required to   remove an electron from a particular subshell  is shown on the horizontal axis, and then the   vertical axis tells us how many of those electrons  are in that subshell. Further to the left   means a greater binding energy which means  electrons that are closer to the nucleus, starting
• 36:00 - 36:30 with the 1s electrons. Then moving to the right  they get lower in energy and farther from the   nucleus. When we see these spectra we should be  able to recognize which peak corresponds to whch   subshell based on its position on the horizontal  axis, and also state how many electrons are in   each subshell based on the height of the peak. We also want to learn about electron affinity.   This is precisely the reverse concept of  ionization energy, it is the energy change
• 36:30 - 37:00 involved with adding an electron to a neutral  atom in the gas phase, thus forming a negatively   charged ion. This process could absorb energy  or release energy, depending on the element, and   a negative electron affinity will mean that the  process is actually favorable for a given element.   Looking at this table, we can see that the  trend is similar to the ionization energy trend,   since the harder it is to remove an electron,  or the higher the effective nuclear charge,
• 37:00 - 37:30 the easier it is to add an electron, and thus a  greater electron affinity. This is why elements   like fluorine and chlorine have very large  electron affinities, as gaining an electron will   endow them with noble gas electron configuration,  which is a very stable situation. So in general,   electron affinity increases going up and  right on the table, with some exceptions.   Noble gases do not follow this trend, as with  a full shell of electrons, it is typically
• 37:30 - 38:00 not favorable to add another electron, so we  discount them when considering this property.  And finally, let’s examine electronegativity.  Electronegativity is a measure of how well   an atom can attract electron density towards  itself, which is measured by looking at the way   electrons are shared in chemical bonds. The more  strongly it can attract electrons, the greater   its electronegativity. Electronegativity will  depend on atomic radius, since a smaller atom with   a greater effective nuclear charge will attract  electrons more strongly, so the electronegativity
• 38:00 - 38:30 trend will be the same as the ionization energy  trend, it will increase going up and right on   the periodic table. Fluorine will have the  greatest electronegativity, and francium will   have the lowest. Again, we will exclude the noble  gases from this trend, as with their full valence   shells, they are not likely to share electrons,  making electronegativity meaningless for those   elements. There is a common point of confusion  that we should make abundantly clear. We must
• 38:30 - 39:00 make the distinction between electronegativity and  electron affinity, because the latter involves an   actual ionization and an associated energy change  that is measurable. The former just describes a   relative calculation of how well an atom attracts  the electrons in a bond, which does not involve   any transformation, and it is listed on an  arbitrary relative scale from zero to four. And
• 39:00 - 39:30 so to put it all together, atomic radius increases  going down and left on the table, while ionization   energy, electron affinity, and electronegativity  all increase going up and right on the table.  And that concludes a review of Unit 1. I’ll  see you over in Unit 2 for more chemistry.