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AP Chemistry > Chapters 7 and 8 > Flashcards

Chapters 7 and 8 Flashcards

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1
Q

The quantum numbers resulted from the…

A

Pauli exclusion principal (two electrons must have opposite spins if they occupy the same orbital)

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2
Q

Nuclear charge is based on the fact that….

A

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

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3
Q

Z eff = ____ - _____

A

Z (atomic #) - S (shielding/# of core electrons)

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4
Q

Effective nuclear charge is the net charge experienced by….

A

a particular electron in an atom resulting from a balance of the attractive forces of the nucleus and the repulsive forces of other electrons

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5
Q

Z eff is the charge experienced by the _____ electrons

A

outermost

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6
Q

High Z eff vs lower Z eff

A

The higher effective nuclear charge has a smaller atomic radius, a higher IE, and a higher electronegativity (periodic trends)

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7
Q

In the hydrogen atom 2s and 2p subshells have the ____ energy, but in atoms with more than two electrons the energies are ____

A

same different

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8
Q

States that lower energy orbitals fill first (predicts arrangement of electrons)

A

Aufbau Principal

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9
Q

Degenerate orbitals (those of the same energy) are filled with electrons until all are half filled before the pairing up of electrons can occur

A

Hund’s Rule

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10
Q

What is the basis of the Hund’s rule?

A

It results in the most stable arrangement of the electrons

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11
Q

When calculating effective nuclear charge (slaters law) it is best to write the electron configuration first and then….

A

group it together by increasing n

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12
Q

4f block of the periodic table

A

lanthanides

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13
Q

5f block of the periodic table

A

actinides

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14
Q

Exceptions when writing electron configurations occure when…

A

you get to d4 and d9 electron configurations

For example:

4s^2 3d^4 becomes 4s^1 3d^5

4s^2 3d^9 becomes 4s^1 3d^10

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15
Q

To form a cation, the highest n electron is removed first for example…

A

Iron (Ar)3d^6 4s^2 becomes Iron (Ar) 3d^6 + 2e-

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16
Q

How do we know the configurations of ions?

A

By their magnetic properties

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17
Q

Paramagnetic

A

has unpaired electrons (attracred to a magnetic field)

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18
Q

Diamagnetic

A

has all paired electrons (not attracted to a metal field)

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19
Q

Ferromagnitism

A

when the spins of unpaired electrons in a cluster of atoms/ions align themselves in the same direction even without a magnetic field

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20
Q

Orbital energies_____as Z* increase

A

drop

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21
Q

distance betweeen two nuclei

A

bond distance (half of this equals the covalent radius, which is usually in picometers)

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22
Q

Ionization energy is always a _____ value

A

positive

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23
Q

Ionization energy is….

A

the energy required to remoced an electron from an atom in the gaseous phase

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24
Q

The _____ to the nucleus an electron is, the higher the ionization energy

A

closer

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25
Q

Electron affinity is….

A

the change in energy (in kJ/mole) of a neutral atom (in the gaseous phase) when an electron is added to the atom to form a negative ion

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26
Q

Noble gase have _____ affinity for electrons

A

no

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27
Q

EA (ΔU) is the energy released when an electron attached to a gas-phase atom to form an anion

Compared to EA, ΔU…

A

Is equal in magnetude, but opposite in sign

EA= -ΔU

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28
Q

Anion>Atom>Cation

Why?

A

Smaller Cation:

Same amount of attractive forces from protons are exerted over less electrons, so the electrons are drawn closer to the nucleus

Larger Anion:

Increased electron-electron repulsion (more electrons than protons)

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29
Q

Isoelectronic

A

Same number of electrons, different number of protons

30
Q

Metallic character

A

How readily an atom can lose an electron

31
Q

Periodic trends from over and up

A

Atmoic radius: decreases

Metallic character: decreases

Electron affinity: increases

IE: Increases

Electronegativity: Increases

32
Q

Electron attachment enthalpy

A

Electron Affinity = -ΔH

33
Q

Covalent bond

A

sharing of valence electrons between atoms (non-metal to non-metal)

also responsible for the atom-atom connections in polyatomic ions

34
Q

Ionic bond

A

Electrostatic in nature, complete transfer of electrons from a metal to a non-metal (cation + anion)

35
Q

Metallic bond

A

attractive forces holding pure metals together, cations are in a “sea” of electrons

36
Q

Core electrons (are or are not?) involved in bonding or chemical reactions

A

Are not

37
Q

Introduced the Lewis Dot Structure

A

Gilbert Newton Lewis

38
Q

Octet rule

A

a chemical rule of thumb that reflects observation that atoms of main-group elements tend to combine in such a way that each atom has eight electrons in its valence shell, giving it the same electronic configuration as a noble gas (stable configuration)

39
Q

Exceptions to the octet rule/don’t obey octet rule

A
  1. More than 4 bonds/8 electrons (expanded octet)
  2. Odd number of electrons
40
Q

When drawing a lewis dot structure which atom is the central atom?

A

The least electronegative

41
Q

When drawing the lewis dot structure for an acid, where does the H+ go and why?

A

It is placed on the outside because acids want to donate the H+ ion, which allows for easy removal

42
Q

Oxoacid

A

An acid containing oxygen

43
Q

Exceptions to Lewis Dot Structures (NASL)

A

N = 8 for all elements except…

H, He =2

Be = 4

B = 6

44
Q

isostructural

A

same structure

45
Q

Formal charge is calculated by…

A

FC = VE - (LP + 1/2 BP)

46
Q

(+) FC…

(-)FC

A

contributes more e- than it receives

gets more e- than it receives

47
Q

A formal charge of 0 means that specific lewis dot structure is…

A

more favored, more stable

48
Q

Resonance structures are not seperate structures they are actually…

A

hybrid/composite structures (not alternating back and forth between the different structures)

49
Q

Resonance _____the energy by distributing electron density over the entire molecule

A

stabilizes/lowers

50
Q

coordinate covalent bond

A

when a bonding pair of e- originates on one of the bonded atoms (just short of fulfilling octet)

51
Q

free radicals have _____ electrons

A

unpaired

52
Q

Atoms from the _____period or __ block and beyond can have an expanded octet

A

3rd

d

53
Q

Bond angles aren’t exact because…

A

the more lone pairs there are, the more pushed together the atoms are, which results in a slightly smaller bond angle

54
Q

Strengths of repulsion

A

lone-lone > lone-bond > bond-bond

55
Q

Molecular geometry vs electron-pair geometry

A

molecular is more general while electron-pair is more specific and takes the lone pairs into account

56
Q

equatorial positions

A

lie in the imaginary sphere around the central atom

(lone pairs prefer to occupy this space because there is more room)

57
Q

axial position

A

north and south poles of the atom

58
Q

Which EPG has no axial or equatorial postions?

A

Octahedral

59
Q

Electronegativity

A

a measure of the tendency of an atom to attract a bonding pair of electrons

60
Q

Electronegativity is measured on the_____scale

A

Pauling

61
Q

Example of a pure covalent bond

(difference of 0 in electronegativities)

A

C-C

62
Q

Positive and negative ends amount to a ____ bond

A

polar

63
Q

Difference of 0.1 to 0.3 in electronegativity

A

non-polar covalent

64
Q

Difference of 0.4 - 1.7

A

Polar covalent

65
Q

Difference of 1.8 and beyond

A

Ionic

66
Q

All homonuclear (composed only of the same element) diatomic molecules are…

A

non-polar

67
Q

dipole moment

(automatically makes a molecule polar)

A

the magnitufe of charges multiplied by the distance (net dipole)

μ = q (charges) * d (distance)

Units of coulomb meters or debyes

68
Q

The greater the difference in electronegativities the more _____the bond

A

polar

69
Q

bond length

A

distance between the nuclei of two bonded atoms

70
Q

energy required to break a covalent bond (in the gas phase)

A

bond dissociation enthalpy

71
Q

ΔHr = …

A

bonds broken - bonds formed

AP Chemistry (3 decks)

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