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AA: Chemistry > Atomic Structure And Bonding > Flashcards

Atomic Structure And Bonding Flashcards

1
Q

Occupied shells tell us the

A

Period number

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2
Q

Number of valence electrons tells us the

A

Group number

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3
Q

What is an electron cloud

A

A region of space where there is a probability of an electron being found

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4
Q

Emission spectra and the atom

A

-When an element is heated, electrons absorb energy and can be ‘promoted’ to higher energy state (excited state)

-When element relapses from excited state to lower energy state (ground state), energy is released as electromagnetic radiation which we see as colour

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5
Q

Bohr model of the atom

A

-electron in a given orbit has a constant energy t/f called ‘energy shells’
-electrons can only occupy fixed energy levels and cannot exist between two levels
-empty orbits have fixed energy levels
-orbits of larger radii have more energy
-there are sub-shells to each major level

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6
Q

If a shell is not occupied, it still exists. True or False?

A

True

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7
Q

How many electrons can s sub-shell hold

A

2 electrons

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8
Q

How many electrons can p sub-shell hold

A

6

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9
Q

How many electrons can d sub-shell hold

A

10

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10
Q

How many electrons can f sub-shell hold

A
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11
Q

What is an orbital?

A

A region of space in which up to two electrons may be located

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12
Q
A
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13
Q

How many orbitals in s sub-shell

A

1

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14
Q

How many orbitals in p sub-shell

A

3

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15
Q

How many orbitals in d-subshell

A

5

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16
Q

How many orbitals in f sub-shell

A

7

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17
Q

S sub-shell

A

-Lowest energy subshell in every electron shell
- one orbital
-spherical shape

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18
Q

P sub-shell

A

-three orbitals
-can hold up to 6 electrons in total

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19
Q

D subshell

A

-contains 5 different orbitals
-can hold up to 10 electrons
-third shell is first to contain d sub-shell

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20
Q

F subshell

A

-start at fourth energy level
-7 different orbitals
-can hold up to 14 electrons

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21
Q

Subshell defintion

A

A specific energy level within an electron shell. Designated s,p,d,f in order of increasing energy

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22
Q

Pauli exclusion principle

A

States only a maximum of two electrons may be found in a given atomic orbital, and that if an orbital is filled, the electrons will have opposite spin

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23
Q

Aufbau principle

A

States that subshells are filled in order from lowest to highest energy and a lower energy subshell will be completely filled before electrons move into higher energy sub-shell

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24
Q

Hunds rule

A

States that electrons in a partially filled subshell will arrange themselves to form maximum number of half-filled orbitals
-singles before doubles

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25
Summarise difference between an atom in its ground state and atom in an excited state
Atom in ground state has all electrons in lowest possible subshells whereas atom in excited state temporarily has one or more electrons in higher energy subshell than lowest possible
26
Copper electron configuration
1s2 2s2 2p6 3s2 3p6 4s1 3d10
27
Chromium electron configuration
1s2 2p2 2p6 3s2 3p6 4s1 3d5
28
What is atomic radius influenced by
-core charge -number of occupied electron shells
29
Core charge defintion
Measure of the net attractive force felt by the valence shell electrons towards the nucleus
30
Calculation core charge
Number of protons- number of electrons in inner shells
31
Trends in atomic radius down a group
Increases
32
Explain why atomic radius increases down a group
-number of occupied energy levels increases -core charge remains constant
33
Trends in atomic radius across a period
Decreases
34
Explain why atomic radius decreases across a period
- core charge increases-> valence electrons are more strongly attracted to the nucleus -number of occupied electron shells remains constant
35
Electronegativity defintion
Strength with which atoms of an element attract electrons when they are chemically combined with another element
36
Electrongeativity will be higher when
-atomic radius is low -core charge is high-> split into nuclear charge and electron shielding
37
Trends in electronegativity down a group
Decreases
38
Explain why electrongeativity decreases down a group
-number of occupied electron shells increase-> atomic radius increases -nuclear charge increases h/w electron shielding also increases so core charge remains the same - further down a group, the further the outermost electron is from nucleus and there are more electrons between outermost electron and nucleus -results in outermost electrons being less attracted towards nucleus
39
Trends in electronegativity across a period
Increases
40
Explain why electrongeativity increases across a period
-number of occupied electron shells stays the same-> atomic radius decreases - nuclear charge increases but number of inner shell electrons stays the same/ so core charge increases -t/f valence electrons become more strongly attracted to nucleus-> inc. electronegativity
41
Electron shielding
Number of inner shell electrons
42
What happens if we provide enough energy to excite an electron beyond all possible electron shells?
-electron will escape the atom entirely-> atom will be ionised
43
Half equations for ionisation
X (g)-> X+ (g) + e-
44
Ionisation energy
Amount of energy required to remove an electron from each of a mole of gaseous atoms
45
Trend in ionisation energy down a group
First Ionisation energy decreases
46
Explain why first ionisation energy decreases down a group
-number of occupied electron shells increases, so atomic radius decreases - t/f nuclear attraction decreases -electron shielding also increases so core charge remains constant -t/f less energy required to remove an electron
47
Trend in first ionisation energy across a period
First ionisation energy increases
48
Explain why ionisation energy increases across a period
-number of occupied electron shells=constant and atomic radius decreases, so valence electrons are closer to nucleus -core charge increases-> valence electrons are more strongly attracted to nucleus -t/f more energy required to remove an electron
49
Elements with low ionisation energy generally have low electronegativity because
Nucleus does not exert a strong attractive force on electrons
50
Elements with high ionisation energies have high electrongeativity because
There is a strong pull exerted on electrons by the nucleus
51
Why do noble gases have high ionisation energy but low electrongeativity
-increased energy to remove electrons because have full outer shell-> CC of 8+
52
Trend across a period for reactivity of metals
Decreases
53
Explain why reactivity of metals decreases across a period
-core charge increases - electron shielding remains constant-> more energy required to remove an electron-> inc. ionisation energy-> reactivity of metal decreases
54
explain why reactivity of metals increases down a group
-core charge stays the same - number of occupied electron shells and electron shielding increases-> decrease in energy required to remove an electron-> reactivity of metals increases
55
Reactivity of non metals
-atoms have greater attractive force between nucleus and valence electrons when they have a low number of electron shells and a high core charge -non metals with greatest attractive force will be most reactive
56
Explain why francium is more reactive than lithium
-core charge stays the same - francium has larger atomic radius and greater electron shielding compared to lithium because has more inner shell electrons. - valence electrons in lithium have a stronger net attraction to nucleus-> require more energy to remove electron in lithium
57
Relative isotopic mass
Mass of that isotope relative to 1/12 of the mass of carbon-12 atom
58
Relative atomic mass
Average mass of all of the isotopes in an element weighted for their relative abundance
59
RAM calculation
(relative abunadance x relative isotopic mass)/ 100
60
What can we use mas spectrometry to measure
-accurate mass for the isotopes of an element -relative abundance of those isotopes

AA: Chemistry (1 decks)

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