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Grade 12 chem > Unit 2 > Flashcards

Unit 2 Flashcards

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1
Q

Planck’s quantum hypothesis

A

Energy is emitted in small bursts or brackets called a quantum of energy

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2
Q

Photon

A

A quantum of light energy

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3
Q

The photo electric effect (Maxwell)

A

Mid 19th C

Light is an electromagnetic wave

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4
Q

The photoelectric effect (Hertz)

A

1887

The energy of light is dependant on the frequency, not the intensity or brightness of the light

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5
Q

The photoelectric effect (Einstein)

A

Electrons can absorb foxed amounts of light energy, these energy pockets are called photons

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6
Q

Bohr model

A

Electrons can be excited to a higher energy level by absorbing a photon and return to a ground state by emitting a photon

“uneven stair case”

The bigger the difference, the more energy the photon has

2n^2

Described the concept of quantized energy levels

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7
Q

Bohr faileurs

A

His model only worked for the hydrogen atom

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8
Q

Wave or particle

A

Electrons have properties of both waves AND particles

This idea helped Schrodinger create his wave equation

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9
Q

Orbital

A

Region around the nucleus where there is high probability of finding an electron. Can only hold 2 electrons with opposite spins

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10
Q

Heisenbergs uncertainty principal

A

It is impossible to simultaneously know the exact position and speed of an electron within an orbit

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11
Q

Quantum numbers

A

The energy of each subshell within a shell are not equal

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12
Q

Pauli exclusion principal

A

No two electrons can have the same 4 quantum values

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13
Q

Aufbau principle

A

Each electron is added to the lowest energy orbital available

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14
Q

Hund’s rule

A

One electron occupies each of several orbitals at the same energy level before a second electron can occupy the same orbital

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15
Q

Principle quantum number (n)

A

Represents the main shell of electrons (1….)

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16
Q

Secondary Quantum number

A

Divides the shell unto subtle that have slightly different energy levels and tells the shape of the subtle (0- N-1)

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17
Q

Magnetic quantum number (ml)

A

Describes the orientation of several subshells (-e - e)

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18
Q

Spin quantum number

A

In each orbital, there is one electron spinning at +1/2 and one spinning at -1/2

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19
Q

Origin of the quantum theory

A

Energy is not continuous

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20
Q

Stability- Hund’s rule

A

The lowest energy configuration for an atom is the one with all the orbitals of an energy level filled with one electron before adding the second. Electrons will choose to be by themselves if possible

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21
Q

Diamagnetic

A

Atoms with no unpaired electrons

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22
Q

Paramagnetism

A

The spin of the unpaired electrons creates a magnetic dipole moment, they act like tiny magnets

An external magnetic field will cause the spins to align

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23
Q

Isoelectronic ions

A

Have the same number of electrons

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24
Q

Anomalies in electron configurations

A

Cr: [Ar] 4s1 3d5

Cu: [Ar] 4s1 3d10

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25
Q

Lewis theory of bonding

A

Atoms are the most stable when they have the same electron configuration (isoelectric) as the noble gasses

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26
Q

Ionic compounds

A

Ionic formation is generally endothermic

Ionic compounds form when the lattice energy is more exothermic

The overall potential energy much be lowered when forming a compound

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27
Q

Square brackets

A

Show that there was a movement of electrons

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28
Q

Bond length

A

The distance between two nuclei in a covalent bond (average is 75pm)

29
Q

Bond energy

A

The energy required to separate two nuclei in a covalent bond (average is 435kJ/mol)

30
Q

Coordinate covalent bond

A

The shared e- come from one atom

31
Q

Super octet

A

Occurs in the d orbital

C,N,O and F can not form super octets

32
Q

Less than an octet

A

Beryllium is most stable with 2 bonds

Boron is most stable with 3 bonds

33
Q

Formal charges

A

Lewis structures with the smallest formal charges are the most stable

34
Q

VSEPR

A

Valance shell electron pair repulsion theory. Uses the idea that valance shell electron pairs will repel each other

35
Q

AX2

A

Linear

36
Q

AX2E

A

Bent

37
Q

AX3

A

Trigonal planar *

Bond angle= 120

38
Q

AX2E

A

Bent

Bond angle= less than 120

39
Q

AX4

A

Tetrahedral

Bond angle= 109.5

40
Q

AX3E

A

Pyramidal

Bond angle of NH3 is 107

41
Q

AX2E2

A

Bent

Bond angle of water is 109.5

42
Q

AXE3

A

Linear

43
Q

AX4E

A

See saw

44
Q

AX3E2

A

T- shaped

45
Q

AX2E3

A

Linear

46
Q

AX6

A

Octahedral

Bond angle= 90

47
Q

AX5E

A

Square pyramidal

In this case, it does not matter which bond you take off

48
Q

Ax4E2

A

Square planar

49
Q

Valance bond theory

A

Covalent bonds form when two orbitals overlap to produce a new combined orbital containing two electrons of opposite spins

50
Q

Hybridization theory

A

Mixing together two of more orbitals on the same atom to form new orbitals that have equal energy

51
Q

Sigma bonds

A

A bond that is formed when two orbitals overlap and the electron density is concentrated along the axis between the nuclei (single bonds)

52
Q

Pi bond

A

A bond that is formed when orbitals are parallel to each other and they overlap to create two regions of electron density on opposite sides of the axis connecting the two nuclei (double and triple bonds)

53
Q

What does the formation of Pi bonds do

A

Prevents the rotation of the molecules along the sigma bonds

Rotation of double or triple bonds require breaking Pi bonds therefore the energy is higher

54
Q

Hybridization- Lone pairs

A

Lone pairs of electrons can be accommodated in hybrid orbitals as well

They have stronger repulsion with bonding pairs so they cause a slight decrease to the bond angles

Bond angle is calculated from sigma to sigma bond

55
Q

sp shape

A

linear

56
Q

sp2

A

trigonal planar

57
Q

sp3

A

tetrahedral

58
Q

sp3d

A

tetrahedral

59
Q

sp3d2

A

octahedral

60
Q

Delocalized molecular orbitals

A

For compounds is might be easier to share electrons by overlapping multiple p orbitals to form a large Pi bond and a delocalized Pi bond

61
Q

Resonance

A

Two possible lewis structures but neither is correct. Real molecule is a mixture of the tw

62
Q

Thompson

A

1897

Positive matter with negative pockets (plumb pudding)

63
Q

Rutherford

A

1911

Gold foil experiment- discovered in the nucleus

Positive charge concentrated in a nucleus with empty, negative space surrounding

64
Q

Chadwick

A

1932

Discovered in the nucleus

65
Q

Ionic crystals

A

Hard, brittle, can conduct electricity when dissolved in water

66
Q

Metallic crystals

A

Metals, shiny, silver, flexible solids that can conduct electrical and thermal energy

Electron sea provides electrostatic glue holding the atom enters together

67
Q

Molecular crystals

A

Non metals

Most are soft crystals with low melting points that can’t conduct electricity

Intermolecular forces hold them together

68
Q

Covalent network crystals

A

Metalloids/ carbon

Hard brittle solids with very high melting points that do not dissolve and can’t conduct electricity

Network of covalent bonds

Grade 12 chem (5 decks)

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