Chemistry 101 Final Unit 1

Question 1

What are elements, compounds and mixtures?

Answer 1

• element: only one kind of atom, cannot be broken down physically or chemically
• compound: 2 or more elements chemically bonded into molecules or extended structures + only broken down into component elements through chemical changes
• mixture: heterogeneous or homogeneous, 2 or more elements physically bonded, can vary in their parts by pass + will retain many of the properties of its components + only broken down through physical changes

Question 2

What is Dalton’s Atomic Theory + the mass laws it accounted for?

Answer 2

• Dalton’s Atomic Theory: matter consists of indivisible atoms,
  atoms cannot be converted to atoms of another element but
  can combine with other types of atoms to form compounds
  
  • incomplete, doesn’t explain how bonds form + new
    experiments suggest an atom could be divisible into
    subatomic particles (electrons + nucleus)
• 3 mass laws
  
  • conservation of mass: total mass is same in chemical in a
    chemical reaction
  • constant (definite) composition: all samples of a given
    compound consist of same fraction (ex. calcium carbonate)
  • multiple proportions: in different compounds of same
    elements, masses of combining elements can be
    expressed in simple ratios (ex. CO + CO2)

Question 3

What did Thomson discover?

Answer 3

1. existence of negatively charged particles + that they’re smaller and lighter than atoms
2. proposed plum pudding model of atom: maintains neutral charge of atom (area of positive charge + negatively charged particles)

Question 4

Atoms can lose or gain electrons to form charged species called?

Answer 4

Ions (atomic number: # of protons = # of electrons)

Question 5

What are ionic compounds?

Answer 5

• name cation then anion
• add -ide to name of anion
• for metals that can form more than one ion, specify the charge by Roman numerals

Question 6

What are polyatomic ions (ionic compound)?

Answer 6

• consists of 2 or >2 atoms joined by covalent bonds + exists as charged unit (ex. NaCN is not Na + C + N, but Na + CN ions)
• # of O increases: hypo-ite, ite, ate, per-ate

Question 7

What are binary compounds (covalent compound)?

Answer 7

• written in order of increasing EN
• add -ide to name of 2nd element
• add numerical prefixes where necessary (mono prefix is not used for 1st element)
• ex. N2O: dinitrogen monoxide

Question 8

What are acids (covalent compound)?

Answer 8

• produce H+(aq) when dissolved in water (formula has at least one H atom usually written first)
• binary acids have no O: hydro-ic acid (ex. HF(aq): hydrofluoric acid)
• oxyacids contain H+O+non-metal, name depends on oxoanion
  
  • -ate: -ic acid, -ite: -ous acid

Question 9

What is spectroscopy?

Answer 9

Study of interaction of matter with radiated energy.

Question 10

What is wavelength (λ), frequency (v), amplitude + speed?

Answer 10

• wavelength: distance between 2 consecutive crests (ex. distance the wave travels during one cycle)
  
  • units: m, cm, nm
• frequency: # of cycles the wave undergoes per second
  
  • units: s^-1: H2
• amplitude: height of cresti related to intensity of radiation
• speed of a wave: distance it moves per unit time
  
  • formula: C: λv

Question 11

Each atom changes it’s energy by an amount, △Eatom…

Answer 11

When it absorbs or emits one proton (one particle of light) whose energy is related to its frequency (not its amplitude).

Question 12

What is threshold frequency?

Answer 12

Electrons break free when it absorbs a photon of enough energy. Electrons cannot be ejected by saving up energy, whose energy isn’t sufficient.

Question 13

What is absence of a time lag?

Answer 13

Some current flows as soon as light of sufficient energy (frequency) strikes the metal surface.

Question 14

What is atomic spectra?

Answer 14

• an element that’s vaporized + excited will emit a series of fire lines (not a continuous spectrum)
• spectral lines are produced when electrons move from one energy to another energy level
• absorption: excitation from lower allowed energy level to higher energy level (absorbs a proton)
• emission: relaxation from higher allowed energy level to lower energy level, often ground level (emits a proton)

Question 15

What are the 3 parts of Bohr’s Model + its limitation?

Answer 15

1. electron occupies stationary states
   
   • ground state: electrons in first orbit (n:1), closest to nucleus + H atom is in its first energy level
   • excited state: electrons is in any orbit farther from nucleus (n>1)
     
     • ex. if n:2, the atom is in the first excited state
   • n:∞: indicates an electron has been completely removed from nucleus
2. electron doesn’t lose energy while in a stationary state
3. electron undergoes a transition from one stationary state to another by absorbing or emitting photons
   
   • absorption process: nf > ni, △E always +
   • emission process: nf < ni, △E always -
4. limitation: fails to predict line spectrum of any other atom + completely for atoms with more than one electron since e-e repulsions, nucleus-e are unaccounted for.
   
   • H-like species: one electron species
     
     • He+: 1 e, 2 p, z:2
     • Li2+: 1 e, 3 p, z:3

Question 16

What are infrared, visible + ultraviolet series?

Answer 16

• infrared: nf:3, when electron drops from outer orbits (ni:4,5,6…) to n:3 orbit
• visible (Balmer) series: nf:2, when electron drops from outer orbits (ni:3,4,5…) to n:2 orbit
• ultraviolet (Lyman) series: nf:1, when electron drops from outer orbits (ni:2,3,4…) to n:1 orbit

Question 17

What does the smallest energy transition correspond to?

Answer 17

Shortest distance on energy plot since smallest △E corresponds to longest λ.

Question 18

What is electron probability density?

Answer 18

A measure of the probability of finding the electron in some tiny volume of the atom. It decreases with distance from nucleus along line.

Question 19

What are the 4 quantum numbers for the state of an electron?

Answer 19

1. principle quantum number (n)
   
   • allowed values: any + integer
   • size of orbital
2. angular momentum quantum number (l)
   
   • allowed values: 0 to n-1
   • shape of orbital
   • 0:s, 1:p, 2:d, 3:f
3. magnetic quantum number (ml)
   
   • allowed values: -l-+l
   • spatial orientation of orbital (e cloud)
   • 1 s-orbital, 3 p-orbitals, 5 d-orbitals, 7 f-orbitals
4. spin quantum number (ms)
   
   • describes spin of the electron (not orbital)
   • half of electrons are attracted by large external magnetic
     field while other half are repelled by it
   • electrons with same spin value: parallel spin
   • electrons with different spin value: opposing spin

Question 20

Orbital shapes

Answer 20

• s-orbitals are spherical, p-orbitals nucleus lies at nodal plane, d-orbitals orientation of nodal planes always lie between orbital lobes

Question 21

What is Pauli’s Exclusion Principle?

Answer 21

No two electrons can have all four quantum numbers alive, each orbital can be occupied by maximum of 2 electrons of opposing spin.
- ex. specify all quantum #s of the 2 electrons in a He atom
He atom: 2 p, 2 e
1. n=1, l=0, ml=0, ms=+1/2
2. n=1, l=0, ml=0, ms=-1/2

Question 22

How do orbital energies differ in one vs many electron atoms?

Answer 22

The outermost electron feels…
- attraction from positive nucleus
- repulsion by other e-e is somewhat screened/shielded from
the nucleus by other electrons

Question 23

What is the effect of nuclear charge (Zeff)?

Answer 23

The higher the nuclear charge, the greater the pull on the electron. Energy is lowered and that the system is more stable.

Question 24

What happens during shielding?

Answer 24

Inner electrons prevent the outer electron from experiencing the full attraction of the nucleus, making its removal somewhat easier than expected.

Question 25

What does orbital energy depend which factors?

Answer 25

1. greater nuclear charge lowers orbital energy
2. e-e repulsions raises orbital energy
3. electrons in outer orbitals (higher n) are shielded from the full nuclear charge, so they have higher energy
4. orbitals with good penetration (having e-density close to nucleus) have lower energy
   
   • penetration: how well the outer electrons are shielded by
     the inner electrons

Question 26

What is Hund’s rule?

Answer 26

When orbitals of equal energy are available, maximize unpaired spins.

Question 27

What’s special about Cr + Cu during electron configurations?

Answer 27

They have “anomalous” configurations attributed to “special stability” of half filled or completely filled 3d subshells.
- ex. Cr: [Ar]4s^2-3d^4 to [Ar]4s^1-3d^5

Question 28

For electron configuration of ions…

Answer 28

1. metals tend to lose valence electrons to form cations
2. transition-metal atoms can form several cations, with the ns electrons removed before the (n-1)d electrons

Question 28

For electron configuration of ions…

Answer 28

1. metals tend to lose valence electrons to form cations
2. transition-metal atoms can form several cations, with the ns electrons removed before the (n-1)d electrons
3. nonmetals tend to gain valence electrons to form anions

Question 29

Difference between paramagnetic and diamagnetic?

Answer 29

• paramagnetic: unpaired electrons
• diamagnetic: all electrons are paired

Question 30

For Zeff…

Answer 30

• increase Zeff as there’s a net + charge: decrease radius (half the size of neutral atom)
• decrease Zeff as there’s a net - charge: increase radius (2x the size of neutral atom)
• decreases down the group, increases across period
• ex. rank in order of increasing size: K+, S2-, Cl-
  K < Cl < S (more negative ion: more e: bigger size)

Question 31

For atomic radius…

Answer 31

• increases down the group (less Zeff cause electrons aren’t held as tight, so the probability that outer electrons spend most of their time from nucleus increases), decreases across period (electrons are held tighter)

Question 32

What is ionization energy and its trend?

Answer 32

• energy that must be added to remove an electron from isolated gaseous atom or ion
  
  • 1st IE: outermost electron (x = x+ + e)
  • 2nd IE: next electron (x = x2+ + e)
  • each successive ionization needs more energy (IE4 > IE3
    >IE2>IE1)
• decreases down the group, increases across period
  
  • anomaly: Be VS B: B has 1 e in 2p orbital (more energy than
    2s) hence removing this e is easier than removing a 2s e
    from Be which has a full subshell)

Question 33

What is electron affinity and its trend?

Answer 33

• the energy change that occurs when a gaseous atom or ion gains an electron
• EA1 is mostly negative (when an electron is added to a neutral atom, energy is released in an exothermic process) where EA2 is always positive (the energy required to add a second electron to an anion)
• decreases down the group, increases across period

Question 34

What is the difference between reducing agent and oxidizing agent?

Answer 34

• reducing: X loses e, X is oxidized by Y, oxidation # of X increases
• oxidizing: X gains e, Y is reduced by X, oxidation # of Y decreases)

Question 35

Are alkali metals good oxidizing or reducing agents?

Answer 35

Reducing agents, since they want to lose an electron so they want to be oxidized.

Question 36

Are halogens good oxidizing or reducing agents?

Answer 36

Oxidizing agents, since they want to gain an electron so they want to be reduced.

Question 37

How can we tell if a substance is likely a good oxidizing or reducing agent?

Answer 37

• good reducing agents: low IE + EN, small e-affinity, low oxidation state
• good oxidizing agents: high IE + EN, large e-affinity, high oxidationg state