Pre-AP Chemistry at BNDS

 Lesson 1

Chemistry is the study of matter and change

Matter has mass and volume (takes up space) - ex. atoms, molecules, cells, tissues, animals, plants

Change comes in two forms: physical and chemical

Physical changes do not change the chemical nature of an object (no chemical bonds are broken or formed)

Chemical changes change the chemical nature of an object (chemical bonds are broken or formed)

Non-matter (which has no mass) also affects changes in matter and plays a role in chemical reactions

Examples of non-matter include light, sound, heat, friendship, pride, and confusion

Lesson 2

All knowledge in chemistry is based on experiment

For scientists to communicate about different experiments, they need a common system of units and measures

The most common units and measures are:

 - Length (meter, m)

 - Mass (kilogram, kg)

 - Time (second, s)

 - Temperature (Kelvin, K)

 - Volume (liter, L)

 - Density (g / cm^3)

 - Speed (m / s)

To convert between different temperature scales [Celsius (C), Fahrenheit (F), Kelvin (K)]:

 - T(K) = T(C) + 273.15

 - T(F) = (9/5)*T(C) + 32 

In communicating very small or very big numbers, scientists use scientific notation to avoid writing lots of extra zeros:

 Ex.  12,340,000 = 1.234 x 10^7 = 1.234 E +7

 Ex. 0.0000789 = 7.89 x 10^-5 = 7.89 E  -5

Scientists also need a way to indicate how precisely they measured a chemical or physical property or change

Ex. the length of a chopstick

Significant figures indicate how precisely a property or change was measured

Important word for counting the number of significant figures in a number:

 - Trailing zero: any zero after the last non-zero number

The number of significant figures in a number ...

 ... with a decimal point equals the total number of non-zero numbers plus all trailing zeros

 ... without a decimal point equals the total number of non-zero numbers

When adding or subtracting two numbers, the total number of places after the decimal should equal the number of places after the decimal in the number with the fewest places after the decimal of the numbers added or subtracted

When multiplying or dividing two numbers, the number of significant figures of the final answer should equal the number of significant figures of the number with the fewest significant figures of the numbers multiplied or divided

Lesson 4

All matter is made of atoms

Two or more atoms combine to make molecules

Matter can be separated into substances and mixtures

Substances have a uniform composition and do not vary from sample to sample

 - Pure elements are made of only one kind of atom (element)

 - Compounds are made of more than one element, but the elements are combined in fixed ratios

Mixtures contain more than one element and may vary from sample to sample

 - Homogeneous mixtures (solutions) are uniform throughout (look the same everywhere in the solution)

 - Heterogenous mixtures are not uniform, some regions look different than others

We can use the physical properties of the parts of mixtures to separate the parts using different separation techniques

 - Extraction: separating A and B in a mixture because A and B are soluble in differnt solvents

 - Decantation: separating a liquid from a precipitate by pouring the liquid off the precipitate

 - Recrystallization: separating A and B by heating A and B in the same solution, then cooling; if A is less soluble, then it will form a crystal precipitate before B, which will stay dissolved in solution

 - Filtration: separating A and B based on the different sizes of A and B molecules using filter paper

Lesson 5

The modern picture of the atom was discovered over many hundreds of years. You know show a few important discoveries.

Conservation of Mass:

 - Discovered by Lavoisier

 - Matter can not be created or destroyed in a chemical reaction

We know that atoms contain electrons, protons, and neutrons. You should know who discovered each and what experiment they used (except for the neutron)

Discovery of the Electron:

 - Mass-to-charge (m/q) ratio of the electron discovered by Thomson using the Cathode Ray Tube

 - Charge of the electron discovered by Millikan using the Oil Drop Experiment

Discovery of the Proton:

 - Discovered by Rutherford by firing alpha particles at gold foil

 - Only 1 / 20,000 alpha particles bounced back!

 - The atom is mostly empty space, with the nucleus having a radius of about 10^-15 m

Discovery of the Neutron:

 - Discovered by Chadwick

Atoms are extremely small: on average, atoms have a radius of 10^-10 m.

Lesson 6

The smallest unit of an element is an atom.

Atoms contain a nucleus of protons and neutrons surrounded by a cloud of electrons.

Chemical symbols allow us to identify an element and count the number of protons, neutrons, and electrons in a particular atom of that element.

Atomic Number (Z) = # of protons in one atom of an element = the same for all atoms of that element

Charge (q) = # of protons - # of electrons in one atom of an element

Atoms with charge are called ions

 - Ions with positive charge are called cations

 - Ions with negative charge are called anions

Atoms with no charge are called neutral (NOT neutron)

The mass of a proton ~ 1 atomic mass unit (amu)

The mass of a neutron ~ 1 atomic mass unit (amu)

The mass of an electrons ~ 1/1800 atomic mass units (amu)

The atomic mass of an atom is the total number of protons and neutrons in that atom (because electrons have so little mass, we do not include their mass in the atomic mass)

Isotopes are atoms with the same number of protons but different numbers of neutrons

The weighted (average) mass of an element is found by taking the sum of the products of each isotope's mass and percent natural abundance (or relative abundance) for each isotope of that element

To find the different isotopes and the relative abundance of different isotopes for an element, we use a mass spectrometer

To use a mass spectrometer, we must first add charge (ionize) an atom by adding or removing electrons

A mass spectrometer separates ions based on the mass-to-charge ratio (m/q) of the ions

Lesson 7

Chemical reactions depend on the interaction of different atoms

Atoms interact by sharing or tranferring electrons

To understand chemical reactions, we need to understand how electrons are organized in atoms

The Bohr model has electrons in circular orbits around the nucleus

These orbits are quantized - electrons can only exist in specific circles around the nucleus, not everywhere

We number different orbits as n = 1, n = 2, n = 3, etc., with a higher n meaning farther from the nucleus

Electrons in different orbits have different energies

Electrons can jump to a higher-energy orbits (farther from the nucleus) by absorbing energy

Electrons can jump to a lower-energy orbits (closer to the nucleus) by emitting energy

Light is one form of energy that can be absorbed or emitted

The energy of light is given by the formulas E = hf = hc / lambda

 - h = Planck's constant = 6.626 x 10^-34 Js

 - f = frequency = units of 1/2

 - c = speed of light = 3 x 10^8 m/s

 - lambda = wavelength = units of length (usually nm)

One way for an electron to jump from a lower-energy orbit to a higher-energy one is to absorb light with energy equal to the difference in energies between the two orbits

 - Ex.  An electron can jump from n = 1 (with Energy = 100) to n = 2 (with Energy = 500) if it absorbs light with an energy equal to E(n = 2) - E(n = 1) = 500 - 100 = 400

One way for an electron to jump from a higher-energy orbit to a lower-energy one is to emit light with energy equal to the difference in energies between the two orbits

- Ex. An electron can jump from n = 2 (with Energy = 500) to n = 1 (with Energy = 100) if it emits light with an energy equal to E(n = 2) - E(n = 1) = 500 - 100 = 400

Since electrons can only exist in specific - quantized - orbits, they can only emit specific - quantized - energies of light, corresponding to specific - quantized - wavelengths of light

The line spectra of different atomic gases clearly demonstrate the quantized nature of energy levels and orbits in an atom

Lesson 8

The Bohr Model is good at predicting the atomic line spectra of atoms (and ions) with only one electron, but it fails for atoms with more than one electron

Instead placing electrons in orbits of a specific radius, the modern model of the atom places them in orbitals that give the probability of finding an electron in a specific volume

To understand the reason electrons are organized in orbitals and to understand the nature of orbitals, we need to understand how and why electrons have both wave and particle properties

We are used to thinking of electrons as particles - little balls of negative charge with a very small mass

Three important ideas tell us why we must think of electrons as waves in atoms:

1) The average radius of many atoms is about 10^-10 m

2) All pieces of matter - big and small - have a wavelength, and the wavelength of electrons is about 10^-10 meters

3) The Uncertainty Principle tells us that we can't know position and velocity of a particle exactly, and the uncertainty in the position of an electron is about 10^-10 m

Since the wavelength of an electron and the uncertainty in the position of an electron are both about the same size as the radius of an atom, we must think of electrons as waves if we want to fully understand how they behave in an atom

To figure out exactly what an 'electron wave' looks like, we need to solve a very complicated problem (Schrodinger Equation)

However, we can get most of the information we need about different 'electron waves' by looking at the quantum numbers of different orbitals

We put electrons in orbitals and electrons can move between orbitals in the Modern Model of the Atom just like we can put electrons in orbits and electrons can move between orbits in the Bohr Model of the Atom

Let's compare Bohr's orbits and modern orbitals

                                                                                                   Bohr Model of the Atom                   Modern Model of the Atom

Electron position is quantized?                                                                Yes                                                    Yes

Electron energies are quantized?                                                            Yes                                                    Yes

Electrons travel around nucleus in ...                                               Orbits (Circles)                 Orbitals (with many different shapes)

Electrons with different energies exist in different ...                      Orbits                                               Orbitals

Electron position described by quantum numbers ...                                  n                                                   n, l, ml, ms

Electrons can jump between different ...                                              Orbits                                               Orbitals

Examples

 

We see that orbits and orbitals are very similar

Putting It Together (Part I):

- Electrons with different energies are in different orbitals.

- Different orbitals have different shapes.

- Electrons with different energies are in different orbitals with different shapes.

Important point:

 - When we say an electron has n = 1 in the Bohr model, we mean that the electron is a specific distance from the nucleus

 - When we say an electron has n = 1 in the Modern model, we mean that an electron has a probability of of being a specific distance from the nucleus

Big idea: Orbitals = Probabilities. Say this 200 times. Say it in your sleep. Ask Mr. Shank 10^9 questions until you know what this means.

Putting It Together (Part II):

 - Electrons in different orbitals have different probabilities of being a specific distance from the nucleus

 - Different orbitals have different shapes

 - The shape of an orbital tells you something about the probability that an electron is a specific distance from the nucleus

We describe how an orbital looks using three quantum numbers:

 - Principle Quantum Number (n) = size of the orbital; possible values are n = 1, 2, 3, 4, ...

 - Angular Momentum Quantum Number (l) = shape of the orbital; for a specific n, possible values are l = 0, 1, 2, ... (n-1)

    Ex. If n = 1, then (n-1) = 0 --> the only possible value of l is 0; if n = 2, then (n - 1) = 1 --> the possible values of l are 0 and 1

    Important notation: We say l = 0 is an s-orbital,  l = 1 is an p-orbital,  l = 2 is an d-orbital,  l = 3 is an f-orbital

 - Magnetic Quantum Number (ml) = orientation of an orbital; for a specific l, possible values are -l to +l

   Ex. If l = 1, possible values of ml are -1, 0, 1; if l = 3, possible values of ml are -3, -2, -1, 0, 1, 2, 3

Spin is a property of an electron - like charge, like mass - that affects how an electron behaves near a magnet

We use one more quantum number to describe the spin of an electron

 - Spin Quantum Number (ms) = spin of an electron; only possible value are ms = +1/2 or ms = -1/2

    Does not depend on any other quantum numbers

    Do not worry about how electrons with different spins act near magnets

 Important Idea: No two electrons can have the same set of four quantum numbers (Pauli Exclusion Principle)

Lesson 9

Chemistry and chemical reactions depend on the transfer or sharing of electrons between atoms.

Remember, the ground state of an atom is its lowest-energy state.

Atoms are in their ground state most of the time.

To understand chemistry and chemical reactions, we need to know how electrons are organized (which orbitals they occupy) in the ground state.

An orbital is a space around the nucleus given by a unique set of (n, l, ml)

    Ex. (2, 1, 0) is an orbital, and (2, 1, -1) is an orbital, but (2, 1) is NOT an orbital.

Golden Rule: In the ground state, electrons occupy orbitals such that the atom has the lowest energy possible.

How to determine the ground-state electron configurations of neutral atoms:

  1) Find the number of protons in the nucleus

  2) The number of electrons in the neutral atom equals the number of protons in the nucleus

  3) Add all the electrons in the atom into orbitals -

         a) lowest-energy orbitals first (Golden Rule)

        b) no more than two electrons per orbital; if two electrons in one orbital, spins must be different (Pauli Exclusion Principle)

        c) if orbitals have the same energy (same n and l), add one electron to each orbital with a different ml before adding a

            second electron to any orbital; first (2l + 1) electrons should have spin in the same direction (Hund's Rule)

        d) remember exceptions for Cr, Cu, Ag, and Au - for these elements, the atoms have lower energy when an electron moves

            from an ns orbital to form a half-filled (5 out of 10 total e-) or fully-filled (10 out of 10) (n-1)d orbital

To determine the electron configurations of ions, first write the electron configuration of the neutral atom, then ...

  1) ... for cations, remove electrons from orbitals with the highest principle quantum number (n) first; if more than one electron has the same n, remove the one with the highest angular momentum quantum number (l) first

  2) ... for anions, add electrons to orbitals in the same order you use for neutral atoms

Lesson 10

Why are electron configurations important?

 - Chemistry IS the sharing and transfer of electrons

 - The configuration of electrons in an atom determines which electrons participate in chemical reactions, and how they participate

Why is the Periodic Table so important?

 - When arranged in order of increasing atomic number, the elements on the table show repeating (periodic) trends in physical and chemical properties.

 - This perodicity comes from

     1) the similarity of the valence electron configurations of elements in each group (column)

     2) the gradual and systematic change in the valence electron configurations of elements in each period (row)

The simplest chemical reaction and chemical bond is the ionic bond:

 - Atom A transfers electrons to atom B, making A positively charged and B negatively charged

 - Since A and B have opposite charges, they attract each other

 - Large systems of postively-charged cations and negatively-charged anions form crystals with interesting properties

Before we understand more complex bonds and chemical reactions, we should understand the formation of ionic solids

Before we understand ionic solids, we should understand why atoms become ions and the size of different ions

Before we understand the size of different ions, we should understand the size of different neutral atoms

So ... let's get to it!

Defining the radius of an atom as the radius within which we have 100% certainty of locating an electron is not practical, since we know that there is always some probability of finding an electron inifinitely far from the nucleus in an atom.

There are many more practical ways to measure the radius of an atom, but one of the simplest ways - the way we'll think about it for now - is to imagine two atoms of the same element coming together to form a molecule. The atomic radius of that element is then half the distance between the nuclei of those two atoms.

From Lesson 10, Slide 17, we see two trends in atomic radius

 -  Atomic radius increases from top to bottom going down a column

 - Atomic radius decreases from left to right going across a row

We need two different ideas to understand these different periodic trends.

Column trend:

 - elements have similar valence electron configurations going down a column

   Ex. H: 1s1, Li: [He] 2s1, Na: [Ne] 3s1, K: [Ar] 4s1,  ... [noble gas] ns1 for elements in Group 1

         He: 1s2, Be: [He] 2s2, Mg: [Ne] 3s2, Ca: [Ar] 4s2, ... [noble gas] ns2 for elements in Group 2

 - we see that the principal quantum number, n, of the outermost valence electron increases by 1 for each element going down a column

 - we know that n is related to the size of an atom, and as n increases, the size of an atom increases

 - so ... it makes sense that for similar valence electron configurations, elements with larger n have a larger atomic radius (i.e. a larger size)

 - summary: atomic radius increases down a column as n of the outermost valence electrons increases down a column

Row trend:

 - all elements going across a row have their outermost electrons in orbitals with the same value of n, so we need a different idea to explain the row trend in periodic radius

 - negatively charged electrons in an atom are attracted to the positively charged nucleus

 - negatively charged electrons in an atom repel the other negatively charged electrons

 - core electrons - closer to the nucleus - repel valence electrons, causing them to 'feel' less of the positive charge of the nucleus

 - the effective nuclear charge, Zeff, is the nuclear charge felt by a valence electron; it is always less than the true charge on the nucleus

 - we calculate the nuclear charge using the formula Zeff = Z - (# core electrons)

 - the Zeff felt by valence electrons in atoms increasess when looking at elements from left-to-right across a row

 - electrons feeling a higher Zeff are pulled closer to the nucleus, so as Zeff increases from left-to-right across a row, the atomic radius decreases

 - summary: atomic radius decreases from left-to-right across as row as the effective nuclear charge felt by valence electrons increases

Lesson 11

 - The simplest interaction between two atoms is the ionic bond:

    * Atom A gives electrons to atom B

    * Atom A becomes positively charged (a cation), and atom B becomes negatively charged (an anion)

    * Atom A and atom B attract each other, since they have opposite charges

 - Atoms in the s-block and the p-block tend to form specific types of ions when making ionic bonds

 - Why do they form specific types? To fill all of their valence orbitals!

 - Atoms that fill their valence orbitals tend to have low energy and be especially stable (i.e, less reactive)

   * Ex. Noble gases, with full valence orbitals, interact weakly with other atoms and form very few molecules

 - Other atoms gain or lose electrons in order to attain full valence orbitals - the number of electrons they need to gain or lose

   determines the specific type of ion they tend to form

    Ex. Elements in Group 1 (alkali metals: Li, Na, K, etc.) have only one valence electron; they tend to lose this electron in order

          to have full valence orbitals, becoming +1 cations

    Ex. Elements in Group 17 (halogens: F, Cl, Br, etc.) only need one electron in order to have completely full valence orbitals, and

          they tend to gain one electron to become -1 anions

 - Atoms in Groups 1, 2, and 13 tend to lose electrons, becoming +1, +2, and +3 cations, respectively

 - Atoms in Groups 15, 16, and 17 thend to gain electrons, becoming -3, -2, and -1 anions, respectively

 - Metal atoms can donate electrons to non-metal atoms, and the metal cations and non-metal anions can combine to form large salt crystals

 - The ionic radius of a cation is always smaller than that of a neutral atom of the same element, since removing one electron reduces the repulsion felt by other electrons, causing the other electrons to move closer to the nucleus

 - The ionic radius of an anion is always bigger than that of a neutral atom of the same element, since adding an electron increases the repulsion felt by other electrons, causing the other electrons to move farther from the nucleus

 - Since some elements tend to gain electron and others tend to lose electrons, trends in ionic radius are a little weird ... don't worry about memorizing them!

 - Atoms with the same number of electrons are called isoelectronic

 - When comparing two isoelectronic atoms or ions, the one with the greater positive charge has the smaller atomic radius; likewise, the one with the greater negative charge has the larger atomic radius.

Lesson 12

 - Atoms transfer electrons in forming ionic bonds: atom A 'donates' one electron to atom B

 - Since atoms A's valence electrons are its outermost electrons, the atoms A's valence electrons are the ones that first 'see' atom B

 - The electron donated by atom A is a valence electron, and it donates its electron into a valence orbital in atom B

 - We have many ways of thinking about and representing chemical reactions

    1) A block of sodium interacting with a cylinder full of chlorine gas to form NaCl

    2) A simple chemical equation: Na + Cl --> NaCl

    3) By showing electrons transferred between atoms using

             i) electron configurations

            ii) box diagrams

            iii) energy-level diagrams

 - Many of these models, however, are either too simple or too complicated; they include too many or too few details about the atoms involved in the chemical reaction to be both useful and helpful in quickly illustrating a chemical reaction

 - Lewis structures are used by chemists to label and summarize the most important interactions between atoms and molecules in a chemical reaction

 - To draw the Lewis structure of a single atom:

    * First, determine the number of valence electrons in s and p orbitals

    * Write the chemical symbol (one- or two-letter abbreviation) of the element to represent the nucleus, core electrons, and any

       valence electrons in d- or f-orbitals

     * Draw the remaining s- and p-electrons as dots around the chemical symbol

            1) For group 2 elements (alkaline earth metals), you may draw the first two electrons paired on one side

           2) For all other groups, first draw up to four unpaired electrons on four different sides of the chemical symbol before

                pairing any two electrons on one side

     * The Lewis structure is complete when all s- and p-electrons have been placed around the chemical symbol

 - For cations (anions), simply remove (add) electrons in the reverse (same) order as you would in drawing electrons around the

    chemical symbol with one fewer (more) proton (who doesn't love parentheses?)

  - The octet rule says that atoms in chemical bonds tend to gain, lose, or share electrons such that they gain a 'full octet' of

     eight valence electrons; Lewis structures allow us to easily see this as four lone pairs of electrons on four different sides of

     an atom

 - For elements in groups 1 - 3, the octet rule basically restates that atoms tend to gain noble-gas electron configurations in

   bonds to other atoms

 - We can represent the formation of ionic compounds by drawing the transfer of electrons from a metal to a non-metal

Lesson 13

 - In an ionic bond, electrons are tranferred from a metal atom to a non-metal atom

 - An ionic compound is made of many ions held together by ionic bonds

 - The smallest representative of an ionic solid is its formula unit

    * A formula unit contains all the ions making up an ionic compound in the correct ratio

    * The formula unit must be overall neutral (no charge)

 - Characteristic properties of ionic compounds include

     * High melting and boiling points due to the strong attraction between cations and anions

     * Tendency to fracture instead of bend when a force is applied to them

     * Good electrical conductivity, since the dissolved ions help transfer electrons through the solution

 Lesson 14

 - Many molecules contain bonds in which electrons are not transferred, but shared

 - Bonds in which electrons are shared are called covalent bonds

 - Covalent bonds form when the electron-nucleus attraction between two atoms balances the electron-electron and nucleus-nucleus repulsion between those atoms

 - We can show a covalent bond by drawing the electrons between the nuclei of two atoms, represented by each atom's chemical symbol

      * Two electrons are shared in a single bond; we can also represent a single bond as one line between two atoms   C - C

     * Four electrons are shared in a double bond; we can also represent a double bond as two lines between two atoms C = C

     * Six electrons are shared in a triple bonds; we can also represent a triple bond as three lines between two atoms C = C

 - We say that single bonds have bond order = 1, double bonds have bond order = 2, and triple bonds have bond order = 3

 - The higher the bond order, the shorter the bond when looking at a bond between two specific atoms

 - Electrons don't have to to be shared equally between atoms in a covalent bond

 - Electronegativity describes the tendency of an atom to attract or 'steal' electrons in a covalent bond

 - The higher the electronegativity, the more an atom tends to attract electrons in a covalent bond

 - Usually, electronegativity increases with increasing Zeff going left-to-right across a period, and electronegativity decreases going from top to bottom down a group as the valence (bonding) electrons are held farther from the nucleus

 - Bonds between elements with the same electronegativity are called non-polar covalent bonds

 - Bonds between elements with intermediate differences in electronegativity are called polar covalent bonds

     * Electrons are pulled more toward one atom than another

     * One atom thus has a small negative charge ('-' pole) while the other has a small positive charge ('+' pole)

 - Bonds between elements with large differences in electronegativity are ionic

     * Since metals tend to have the smallest electronegativities and non-metals the largest, ionic bonds tend to form between

       metals and non-metals

 - Important point: in determining the polarity of a bond, you must compare the differences in electronegativity, not absolute electronegativities

Lesson 15

 - Lewis structures show us:

     * how many electrons are in bonds between two atoms

     * between which atoms those bonds form

     * how many electrons are found 'isolated' on single atoms

     * on which atoms those electrons rest

 - Many molecules contain more than two atoms

 - Our goal is to draw Lewis structures that best represent how the nuclei and electrons in the molecules are organized in reality

   (in other words, in the lab, in the air, in our body, food, and medicine)

 - To draw Lewis structures for these bigger molecules, it helps to follow a few simple rules

     1. Start with a formula of the molecule, listing how many atoms of each element are in the molecule

     2. Except for hydrogen, place the least electronegative atom in the center

     3. Draw all other atoms around the central atom, with hydrogen and any halogen atoms on the outside

     4. Count the total number of valence electrons in the molecule; this is equal to the sum of the valence electrons of all the atoms in the molecule.

     5. Place single covalent bonds between each of the atoms.

     6. Subtracting two electrons for each single covalent bond, determine how many valence electrons remain.

     7. Place the remaining electrons in the molecule in order to satisfy the octet rule on every atom, starting with the most electronegative atoms (in some cases, you will not be able to make a full octet on every atom)

     8. Sometimes, it may be necessary to make double or triple bonds between two atoms in order to satistfy the octet rule on all atoms. Make double and triple bonds by moving two lone pair electrons from one atom into a bond between that atom and another atom.

After satisfying rules 1 - 8, you may find yourself with two or more possible Lewis structures. To choose the structure that most likely matches the shape and bonding of the real molecule, we need two more ideas: formal charge and resonance.

 - The formal charge on an atom in a molecule can be calculated as: formal charge = valence e- - (2*lone pairs + 1*bonds)

     * An easier way is to 'cut' all the bonds surrounding an atom in half; count all the electrons around an atom by

          i) counting one electron for each bond (one half of the electrons in the 'cut' bonds)

          ii) counting each electron not in a bond (so a lone pair of electrons counts as two electrons)

      * Then, substract the total number of electrons around the atom from the number of valence electrons in the neutral atom

Important Point: In trying to satisfy the octet rule on an atom, we count two electrons in each bond. In calculating the formal charge on an atom, we count only one electron in each bond.

     9. Choose the Lewis structure that has

          a) the fewest number of non-zero formal charges

          b) formal charges with the lowest absolute value (i.e., better to have (0) X--Y(+1) than (+2) X--Y(-1)).

Finally, for some molecules we can draw Lewis structures that are almost exactly the same except for the placement of their double bonds and triple bonds; Examples: NO3 (1-) and  H3CCOO(1-) and CO3(2-)

 - By measuring the bond lengths in each molecule, we see that each of the molecules above are not best represented by one Lewis structure; for example, all N-O bonds in NO3(1-) have the same length, which is not what we would expect based on a single Lewis structure that predicts one short double bond and two longer single bonds

 - To explain the identical bond lengths, we need to draw, for example, three different Lewis structures for NO3(1-), placing the double bond between nitrogen and a different oxygen in each structure

 - We call these different structures resonance structures, because the lone pair electrons on one atom and the electrons in the double bond between two other atoms move back-and-forth, or resonate, between all the atoms involved

 - The 'true' Lewis structure of a molecule is a combination of all of its resonance structures

     10. Draw all equivalent resonance structures for a molecule with double-headed arrows in between each one <-->

Lesson 16

 - Q: Why are we drawing all these Lewis structures? A: Because we want to know which atoms are bonded to which other atoms in a molecule, and we want to know how electrons are organized in molecules.

 - Q: Why do we care about how atoms are bonded in molecules? Why do we care how electrons are organized:

   A: Because we want to explain chemical reactions, and in chemical reactions atoms re-organize their electrons to break bonds and make bonds

 - To fully understand how molecules make and break bonds in chemical reactions, we need to learn about two things:

     1) the geometry (shapes) of molecules

     2) the way electrons are distributed in molecules

 - We need to connect two big ideas to understand the shapes of molecules:

     1) In molecules, there are electrons in bonds, and there are electrons on isolated atoms (lone pairs)

     2) Electrons repel other electrons

 - Valence Shell Electron Pair Repulsion (VSEPR) Theory predicts that molecules take shapes that minimized electron repulsion

 - Consider water (H2O); we see in its Lewis structure that there are two O-H bonds and two lone pairs of electrons

 - This means three different types of interactions between electrons:

     1) Lone pair - lone pair: interactions between the two sets of lone pair electrons

    2) Lone pair - bond pair: interactions between the electrons in one lone pair and the electrons in one bond

    3) Bond pair - bond pair: interactions between the electrons in one bond and the electrons in another bond

 - Since electrons in bonds are 'localized', they occupy a smaller region of space and 'push' other electrons away from them more strongly; this means we can order the repulsion of electrons interactions as:

                                  (most repulsive)  lone pair-lone pair < lone pair-base pair < base pair-base pair < (least repulsive)

 - Using these ideas, we can predict the geometries of molecules with 1, 2, 3, 4, 5, 6, ... groups of electrons around the central atom

Important Point: we count all electrons in between two atoms as one electron group; in other words, two electrons in a single-bond are ONE electron group; four electrons in a double-bond are ONE electron group, six electrons in a triple bond are ONE electron group, etc.

 - You do not need to memorize the geometries of central atoms with different numbers of electron groups for this class, but you should be able to understand and explain them

 - Elements have different electronegativities, which makes bonds between atoms of different elements polar

 - The bigger the difference in electronegativity between two elements, the larger the polarity of the bond between them

 - We can draw the polarity of a bond using an arrow showing the bond dipole; the bigger the arrow, the larger the polarity, the larger the difference in polarity

 - We can add all the bond dipoles in a molecule to find a molecular dipole

 - Molecules with a molecular dipole have negatively charged ends and positively charged ends

Lesson 17

 - In ionic bonds, electrons are transferred; in covalent bonds electrons are shared

 - In both ionic and covalent bonds, electrons are usually either transferred or shared between only two atoms

 - Exceptions include resonance structures, in which electrons resonate between more than two atoms

 - In metallic bonding, all the electrons are shared between all of the cationic nuclei

 - We call the cloud of electrons around the nuclei an 'electron sea'

  - Many properties of metals come from the ability of their nuclei to 'swim' in the electron sea

     * Malleability - the ability to flatten metals into thin sheets

     * Ductility - the ability to pull metals into thin wires

     * Moderately high melting points and very high boiling points

     * High thermal conductivity - the ability to transfer heat

 - Other properties of metals come from the ability of electrons to swim freely around the nuclei

     * High electrical conductivity - the ability to conduct an electric current

 - Alloys are solutions or mixtures of different elements that have metallic properties

 - Alloys are useful because they combine the properties of different elements to make materials that are extra hard, ductile, resistant to corrosion, etc.

Lesson 18

 - We've studied the interactions of ...

   ... electrons and nucleons (protons and neutrons) to make atoms

   ... two atoms to make bonds

   ... two or more bonds to make molecules

 - Next up: studying how two or more molecules interact to make different phases

 - Intramolecular forces (bonds) act between atoms within a molecule

 - Intermolecular forces act between molecules

 - Intermolecular forces can generally be grouped into dipole-dipole interactions, hydrogen bonding, and London dispersion forces

 - Dipole-Dipole Interactions

     * Polar molecules possess a net molecular dipole; this means that one part of the molecule is partially negatively charged

        another part is partially positively charged

     * In a solution of polar molecules, the molecules tend to orient themselves so that the partially positive parts one molecule are

        close to the partially negative parts of another molecule

 - Hydrogen Bonding

     * Hydrogen bonding is a special type of dipole-dipole interaction

     * When covalently bonded to N, O, F, or Cl, hydrogen has a partial positive charge and N, O, F, or Cl has a partial negative

        charge

     * Hydrogen bonding occurs when a hydrogen covalently bound to one of N, O, F, or Cl is also attracted to a N, O, F, or Cl on

       another molecule (or farther away on the same molecule)

     * We can depict this generally as

                                            Covalent Bond   Hydrogen Bond

                                                      |                         |

                                                     V                        V

 

                  |------- >    X --- H - - - - Y  < -----|  

 (X = O, N, F, Cl)          |--------- | ---------| |--------------- | ---------------|         (Y = O, N, F, Cl)

                                        Same Molecule        Different Molecule                              

                                                             (or farther away on same molecule)

 - London Dispersion Forces

     * Intermolecular forces also exist between non-polar molecules

     * Electrons move randomly around all molecules

     * By chance, at times there will be more electrons on one side of a molecule than another, producing an instantaneous dipole

     * The formation of an instantaneous dipole in one molecule can induce an instantaneous dipole in another molecule because the

        electrons on the second molecule are attracted to the positive end of the first

     * Instantaneous and induced dipoles are called London dispersion (or van der Waals) forces

     * London dispersion forces exist between all molecules

     * The number of electrons, size, and shape of a molecule all influence its London dispersion interactions

     * The more easily electrons can be pushed (polarized) into a dipole, the larger a molecule's London dispersion interactions

     * Elements have increasing polarizability (and thus London dispersion forces) as you go down a group

 - In general, intermolecular forces can be listed by increasing strength as: London dispersion < dipole-dipole < hydrogen bonding

 - However, it is possible that molecules with large London dispersion forces can experience stronger net intermolecular forces than molecules with weak dipole-dipole and hydrogen bonding 

 - Intermolecular forces are helpful in explaining trends in melting and boiling points

Lesson 19

 - Before we can talk about the phases and matter and chemical reactions, we need to learn an important concept: the mole

 - Everything we learn in Chemistry had first to be proven by experiment

 - Many, probably most scientists conduct experiments with mg, g, or kg of chemicals, BUT they want to understand the results of their experiments at an atomic or molecular level

 - In other words, we want to learn something about what we can't see through materials and reactions we can see

 - The mole is the bridge between the microscopic world - too small to see - and the macroscopic world - big enough to see.

 - One mole of a substance is 6.002 x 10^(23) particles of that substance. We call 6.002 x 10^(23) Avogadro's Number.

 - We can have one mole of anything ... carbon atoms, water molecules, skin cells, pencils, people, anything!

 - The mole is useful because ...

   ... the mass of one atom of an element in atomic mass units equals the mass of one mole of atoms of that element in grams

      Ex. The mass of one C atom is 12 amu. The mass of one mole of C atoms in 12 g.

   ... the mass of one molecule in atomic mass units equals the mass of one mole of that molecule in grams

      Ex. The mass of H2O is 18 amu. The mass of one mole of H2O molecules is 18 g.

   ... the mass of something in atomic mass units equals the mass of one mole of that something in grams

      Ex. The mass of one pencil is 6.022 x 10^(25) atomic mass units. The mass of one mole of pencils is 100 g.

The mass of one mole of something is its molar mass.

      Ex. H2O has a molar mass of 18 g/mol.

We can read atomic masses on the Periodic Table two ways:

   - as the mass of one atom of an element in amu

   - as the mass of one mole of atoms of an element in grams

The molar mass of a molecule is the sum of the molar masses of the atoms in the molecule.

Mass is conserved in a chemical reaction (remember Lavoisier!)

Atoms will not gain or lose protons in most of the reactions we'll study. (We will look at nuclear chemistry next semester.)

This means that in order to conserve mass in a chemical reaction (remember Lavoisier!), the number of atoms of each element must be equal in the reactants and products.

We place stoichiometric coeffcients in front of reactants and products in order to conserve mass.

     - We call this balancing a chemical equation and the reaction with the correct coefficients a balanced chemical equation

We can read a chemical equation two ways - microscopically and macroscopically. For example, consider the reaction

2Na + Cl2 --> 2NaCl

Microscopically, we read this reaction as:

'Two atoms of sodium react with one molecule of chlorine to produce two formula units of NaCl.'

Macroscopically, we read the reaction as:

'Two moles of sodium react with one mole of chlorine to produce two moles of NaCl.'

The mole (Avogadro's Number) is so important because it allows us to connect these two worlds - big and small - so easily.

Lesson 20

In addition to the mole and molar mass, chemists quantify matter and chemical reactions in other ways

The percent composition of a compound is the percent, by mass, of each element in a molecule of that compound:

% composition (X) in Y = mass (X) in Y/ mass (Y)

% composition (X) in Y = [(number atoms (X) in Y)x(molar mass (X)]/molar mass(Y)

For example, the percent composition of H in H2O is

% (H) in H2O = [(2 atoms H / molecule H2O)x(1.0079 g / mole H)] / (18.0158 g / mol H2O)

% (H) in H2O = 2x1.0079/18.0158

% (H) in H2O = 0.1119

% (H) in H2O = 11.19 %

Note: a careful student will notice that the units in the above calculation didn't cancel. Did we make a mistake? Why not?

Combustion - burning in oxygen - is familiar to everyone that has ever lit something on fire

In a general combustion reaction, a sample is burned in oxygen, producing CO2, H2O, and inorganic products

Hydrocarbons are compounds that contain only carbon, hydrogen and oxygen (CxHyOz)

In combustion reactions for hydrocarbons