AP Chem Unit 1

Unit 1.1 Atomic Structure

Atomic Theory

• compound consists of atoms of two or more elements

• particles attracted by positive charges are repelled by negative charges

  • like charges repel

• particles are less mass than atoms, indistinguishable no matter source, subatomic consists of all atoms

• 1.6 e^-19 is fundamental charge of electron

• electron mass: 1.602 e^-19 C = 9.107 e^-31 kg

• neutron mass = proton mass

Atomic Structure

• nucleus mass = electron mass

• electrons are majority of atoms volume

• unified atomic mass (u): 1/12 mass single carbon-12 isotope = Da (Dalton) = amu (atomic mass unit)

• proton and neutron are 1 amu and

  • proton +1 charge

  • neutrons neutral charge

• electron 0.00055 amu and -1 charge

• # of electrons = # of protons

  • remaining are neutrons

• unequal particles is an ion

• negative charge in anion

• positive charge is cation

Unit 1.2 Mass Spectroscopy

• mass spectroscopy is an analytical tool useful for measuring the mass-to-charge ratio (m/z) of the molecules in the sample

• measurements can be used to calculate the exact molecular weight of the sample and identify unknown compounds using molecular mass

Mass spec Key Ideas

• isotopes are separated based on their mass-to-charge ratio

• relative amounts of each isotope are recorded

Chemical Formulas:

• average atomic mass can be calcuated by

  • AAM = (mass)(relative abundance decimal) + …

• Structural formula does not equal molecular formula

  • 2H does not equal H₂

• diatomics: Hydrogen, Oxygen, Bromine, Nitrogen, Chlorine, Flourine, Iodine

• Molecular formula always whole number multiple of empiracal formula

• same atoms can be arragend differently

  • isomers

  • same chemical formula but different structures

Atomic Masses:

• mass spectrometer converts molecules to ions so they can be moved/manipulated

• 3 essential mass spectrometer functions:

  • ion source (small sample mixed)

  • mass analyzer ( ions separate because mass and change)

  • detector (separate ions measured and displayed)

• molecular ions left after high energy electrons collide with molecule and ionizes it (bonding/nonbonding)

  • fragment ions are residual

• ionizer is vacuum that pulls electrons from sample

• Mass analyzer is electric field that is negatively charged

  • magnet to bend ions

• detector is electron multiplier (spawns more as they hit it) (amplifier)

• calibration sends ions through

• High intensity = peak = lots of ions at weight

• each isotope has different mass

• mass spectroscopy finds abundance of isotopes

Formula Mass and the Mole Concept

• for covalent substances, formula mass = molecular mass

• ionic compound formula does not represent composition of discrete molecules, so not molecular mass

• adiditional/missing electrons/charges can be ignored in mass

• mole of subtance is amount where there are 6.022 e²³ desicrete atoms/molecules

  • 1 mole is same regardless of element, but masses still differ

• molar mass = mass (g) of 1 mol = atomic formula mass

Stoichiometry Review

• mole ratio based on coefficients

• of balanced equation

• 6.022 e²³ avogrado’s number

• 1 mol = 22.4 L

  • 1000 mL = 1 L

• limiting reactant is the reactant present in the smallest stoichiometric amount

  • reactant you will run out of first

1.5 Atomic Structure & Electron Configuration

Electron Configuration

• effective nuclear charge slightly increases as you go down group and dow period

• atomic radius decreases as you go across period and up group

• ionization energy goes down as you go down group and down period

• electronegativity goes up as you go down and across period

• Coloumb’s law - F = (q1q2)/r²

  • r = atomic radius

  • q = charge

• electrons thought to be in shells/sublevels

• inner electrons are core electrons

• outer electrons are valence electrons

• electron configuration: Aufbaus’s Principle

  • an electron occupies orbitals in order low to high energy

• ionization energy (move electron from shells) estimated from Coloumb’s law

  • relative to molecules

• electromagnetic forces (repulsion/attraction) determined by charge being positive or negative and distance between charges

• 90% probability of finding electrons in orbitals (s)

  • maximum 2 electrons occupy the same orbital n=1 (first shell), n=2 (second shell), etc

  • orbital in each shell, 1s 2s, 3s, etc

• Pauli’s Exclusion Principle: no two electrons can have same set of quantum numbers and can’t be same state

  • opposite spins

• Hund’s Rule of Maximum Multiplicity: every orbital in sublevel is singly occupied before any orbital is doubly occupied, all single same spin

Forces in Atom

• outer/valence electrons are:

  • attracted toward positive nucleus

  • repelled by inner electrons

• the balance of these forces leads to binding energy of electron

  • how tightly it’s held orbiting the nucleus

• attractive force between charged particles increases with:

  • increased magnitude of charge

  • decreased distance between the particles

  • forces happen between each charged particle

• protons in the nuclus will attract the electrons

• electrons will repel each other

  • a single orbital can only hold two electrons

  • inner electrons also push other electrons farther away

• different energy levels are called shells

Ground and Excited States

• electrons are generally in lowest energy level

  • if energy is put into an atom, electrons become excited

    • they will leap to higher energy level

1.6: Photoelectron Spectroscopy

• electron’s are misled by x-ray of equal energy

• electrons with more kinetic energy are higher in orbitals, less ionization energy for in orbitals

• KE_proton = IE + KE_electron

• Coloumbic forces what holds the electron in Hydrogen atom

• PES is an experiment technique that measures the relative energy of electrons in atoms or molecules

  • uses photoionization to eject electrons from sample using high energy EM radiation (UV or x-rays)

• PES graphs show relative number of electrons and binding energy (how much energy needed to remove that electron from the atoms)

  • electrons with high binding energy are most strongly attracted to the nuclei (and closes)

  • can read electron configurations from PES graphs

• ionization energy is the minimum amount of energy required to remove the most loosely bound electron of an isolated neutral gaseous atom or molecule

• binding energy is the least energy required to remove a particle from a system of particles

1.7 Advanced Periodic Table Trends

• as atomic number increases, atomic radius decreases because nuclear attraction is stronger and pulls orbitals closer

• electronegativity = amount by which atoms attract electrons

• ionization energy = the energy required to remove one or more electrons

• ion radius = the size of the ion, from the nucleus to the valence shells

• atomic radius = the size of the atoms, from the nucleus to the valence shells

• effective nuclear charge (Z_eff) is the pulling force that a valence electron actually feels

  • Z_eff = number of protons - number of core electrons (all not valence)

• nuclear charge affects properties of the atom like electronegativity, ionization energy and atomic radius

• elements in same representative groups tend to form the same charge ions

  • Group 1A tends to bee +1 charge

• atoms will gain or lose electrons to have a full valence shell electron configuration

• cations lose their electrons and are more positive charged

• anions gain electrons and are more negatively charged