AP Chem Unit 1

Elements

• each one is made up of a unique kind of atom

• substances that cannot be decomposed into simpler substances by chemical methods

• 7 diatomic elements in pure form (H2, N2, O2, F2, Cl2, Br2, I2)

• All of the rest are monatomic

Compound

• made of 2+ kinds of elements

• 2+ capital letters

• can be decomposed to form simpler substances

• compounds have definite compositions (the relative number of atoms in each element is what makes the compound what it is… ex. H20 and H202 are different compounds)

Moles:)

• amount of a substance that contains 6.022 × 10^23 units (particles)

• kind of like a dozen but a different number - its a standard counting unit

• molar mass = atomic mass in g

Converting between moles, molar mass, and mass

mass = moles x molar mass

mol = mass/molar mass

mol = # particles/avogadro’s number

#particles = mol x avogadro’s number

Determining mass of an atom

Mass number (A) = number of protons + number of neutrons (because electrons determine volume, but are too small to determine mass)

Ions and ionic compounds

Created by adding or removing electrons

• cations (positive charge) you get by removing an electron

• anions (negativ charge) you get by adding an electron

Isotopes & Average Atomic Mass

Atoms of the same element (same number of protons) but different masses (because different number of neturons)

• average atomic masses exist because of isotopes existences, weighted by their relative natural abundances

• average atomic weight = isotope mass x fractional natural abundance

ex. Boron is 19.9% B10 and 80.1% B11, so its atomic weight is 10.80 amu

Mass Spectrometry

Technique for measuring the molecular weight and determining the molecular formula of an organic compound

• number of peaks show the number of isotopes and the height of the peaks show their percent abundance

• technique uses mass separation and a magnet (irrelevent)

Percent Composition

% of the mass of a compound that comes from each fo the elements in the compound

% Element = (number of atoms x atomic weight)/(avg atomic weight of the compound) x 100

Empirical Formulas

Give the lowest whole number ration of atoms of each element in a compound

• we can use the percent by weight of each element to derive the empirical formula

Derive it by

1. find percent compositions

2. find masses

3. calculate number of moles of each one based on molar masses

4. divide by the smallest number of moles to get ONE whole number (ex. if there are 7.5 moles of B and 18.75 moles of H, divide both of those numbers by 7.5 to get 1 mole of B and 2.5 moles of H)

5. Multiply by an integer (the same integer) to get whole numbers (so our final formula would be B2H5)

Molecular formulas

Give the exact number of atoms of each element in a compound

• (if we know the molecular formula, we can determine the empirical, but the opposite is not true)

Hydrates

• a salt that when crystalized from an aqueous solution incorporates a fixed amount of water in a crystalline matrix

• the number of moles of water present per ONE mole of anhydrous salt is usually some simple whole number (ex. salt x 5 H2O)

How to do hydrate analysis

1. find all the masses (of the salt, the hydrate, and the water)

2. find all the moles

3. determine formulas

Substances vs. Mixtures

substances are “pure” - they have distinct properties and a composition that does not vary from sample to sample

• can be made up of elements and/or compounds

mixtures are 2+ substances, not chemically combined

• exhibit the properties that make them up

• can be homogenous or a solution (ex apple juice) or heterogenous (ex soil)

• can be separated into pure substances using physical methods

• composition can vary from sample to sample

Distillation

• Uses differences in boiling points of substances to separate a homogenous mixture

Column Chromatography

• separates substances on the basis of differences in the ability of substances to adhere to the solid surface (usually silica gel)

Thin Layer/Paper Chromatography

• separates substances on the basis of differences in the ability of substances to adhere to the solid surface, in this case, dyes to paper

3 Primary Characteristics of Waves (and their relation)

• Wavelength (the distance between two consecutive peaks or troughs)

• Frequency (the number of wavs PER SECOND that pass a given point in space

• Speed of light = 2.998 × 10^8 m/s (constant) in a vacuum

speed of light = wavelength x frequency

• because all electromagnetic radiation travels at the same velocity

The Photoelectric Effect

• energy comes in packets called quanta that can be emitted or absorbed, and the energy of a single quantum is E = hv

• energy is proportional to frequency in that relationship, where h = planck’s constant

• the ability of light to cause chemical change is dependent on frequency (like the difference between UV and visible light)

Bohr Model

• electrons are in definite orbits at a dixed distance from the nucleus, and these are also called energy levels

• further away from the nucleus = more energy (E = hv)

• an atom with electrons in the lowest possible energy level (closest to the nucleus) is said to be in the ground state

• distance between energy levels gets smaller when energy levels are fruther away from the nucleus

Principal Quantum Number (n)

• when electrons move from lower to higher energy levels, they absorb energy. When they come back to their original (ground) state, they emit that energy.

• Principal quantum number is the distance from the nucleus (also dependent on size of orbital, because larger orbitals are farther away from the nucleus)

• Also an indication of energy level

• As n increases, energy also increases

• known as “shell” and has integer values.

How do energy and electrons changing levels relate?

• Electrons can only exist in certain discrete energy levels

• Energy is involved in the transition of an electron from one energy level to another

• Electrons are treated as both particles and waves

• The nucleus is surrounded by orbitals, and they are like clouds showing the probable regions of electrons (you cannot determine the exact location of an electron)

Heisenberg Uncertainty

• psi² gives the electron density, or porbability, of wheree an electron is likely to be at any time

• Heisenberg showed that teh momenutm and position of an electron cannot be known at the same time.

Sublevels

• each energy level is divided into sublevels, which describe the shapes of orbitals

• the number of sublevels within the energy level is equal to the levels principal quantum number

• the sublevels have specific names (s,p,d,f) in order of increasing energy

S orbitals

• spherical in shape

• radius of the sphere increases with the energy level

• each s orbital holds 2 electrons (at maximum)

P orbitals

• have a dumbbell shape

• each have two lobes with a node in between them (and electron density happens in the middle of the dumbell, near the node)

• p orbitals have 6 electrons at maximum

D orbitals

• abstract lobe shapes (usually four lobes)

• d sublevel holds max 10 electrons

• can only happen when n is greater than or equal to 3

F orbitals

• very complicated shapes

• holds 14 electrons at max

• only happens when n is greater than or equal to 4

Electron Configurations

• describes how electrons are distributed in an atom

• the most stable organization is the ground state

ex. 4p^5 denotes that n = 4, p is the type of orbital, and 5 is the number of electrons in that orbital

Orbital Diagrams

Each box in the diagram represents one orbital (group of 2)

Half arrows represent the electrons

The direction of the arrow represents the relative spin on the electron (the first one in each orbital always points upwards, and then the second round of filling them points downwards)

Orbital Filling Rules

1. The lowest energy level fills first (draw all 4 types of sublevels out in a grid and then make diagonal lines to determine order)

2. All orbitals of equal energy (ex all the orbitals in a 2p sublevel) get at least one electron before any of the boxes get 2

3. Electrons in the same orbital must have opposite spins

Orbitals in Periodic Table

• we fill them in order of increasing energy

• different blocks on the periodic table correspond to different types of orbitals

• columns are called groups, rows are called periods

Condensed Electron Configurations

If elements are really far down (so they have a bunch of filled inner cells), we write a shortened version of an electron configuration using brackets around a noble gas symbol and listing only the valence electrons (ex. Sodium is [Ne] 3s^1)

Anomalies when filling shells

When there are enough electrons to half fill s and d orbitals on a given row, you try to fill the d orbital halfway or fully rather than fully filling the S orbital

ex. Cr has [Ar] 4s^1 3d^5 instead of 4s² and 3s^4 because that makes one electron in each “box of the d orbital”

Cu has 4s^1 and 3d^10 instead of 4s² and 3d^9

Ion configurations

• For anions, add 1 or more electrons to the subshell of highest energy

• For cations, remove 1 or more electrons to the subshell of highest energy

• for transition metals, remove the s shell electrons before removing d or f because they are less stable

• valence electrons determine an elements reactivity (ex. noble gases are not reactive) so elements in the same group (column) have similar properties because they have the same valences

Coulomb’s Law

Electrons are both attracted to the nucleus and repelled by other electrons.

• as the number of protons increases, the attraction increases

• as distance increases, attraction decreases (electrons on higher shells are less attracted to the nucleus)

• Valence electrons are repelled by inner shell electrons

• the forces an electron experiences are dependent on both factors

Effective Nuclear Charge (Zeff)

• as more electrons are added, the inner layers of electrons shield the outer electrons from the nucleus (electrons on the same level don’t shield each other)

• Outer shell electrons experience a lower net attraction to the nucleus

It is found by subtracting the atomic number from the shielding constant (approximately the number of inner shell or non-valence electrons). ex. Na has a zEff of 11-10 = 1

*****STRAIGHT RIGHT ARROW FOR THE TREND*****

Atomic Radius on the periodic table

Size decreases across a period because of an increased Zeff

Size increases down a group because there are more shells

Ionic Radius

• atoms form ions to achieve more stable electron configurations

• elements in a same main group form ions of the same charge (ex F and Cl both become F- and Cl-)

Cations get smaller, because they lose electrons (so repulsion decreases)

Anions get larger, because they gain electrons (repulsion increases)

Isoelectronic Ions

Ions of different elements that have the same number of electrons

• ions with more protons are smaller because they are more attracted to the nucleus

Ionization Energy

• Energy required to remove an electron from an atom in the gas phase

• First ionization energy is lower when there is less attraction between valence electrons and the nucleus, and its higher when that attraction is stronger

  • increases across a period

  • decreases down a group

• First ionization is much lower than second ionization, and when you are dropping a shell, that ionization energy is MUCH larger

*****UPWARDS DIAGNONAL TREND*****

Electron Affinity

• how bad something wants to add an electron

• it releases energy (is exothermic), so the value is negative

  • when an atom wants an electron more badly, the EA is more negative

*****UPWARDS DIAGNONAL TREND*****

exceptions are when they have to open a whole new sublevel or when the orbitals each have one electron and adding a second would make more repulsion