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Year 11 Chemistry > Term 1 > Flashcards

Term 1 Flashcards

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1
Q

elements

A

types of atoms differing from each other by the number of subatomic particles

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2
Q

the Bohr model

A

an atom is a dense nucleus surrounded by a cloud of negatively charged electrons in distinct energy levels
> the atom is held together by electrostatic forces of attraction between the protons and electrons

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3
Q

subatomic particles

A

the particles within and surrounding the nucleus

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4
Q

electrons

A

negatively charged particles that travel at certain energy levels (but can jump)

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5
Q

similarities between electrons in same shell

A

similar distance from nucleus, and have similar energy

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6
Q

which electrons have the lowest energy?

A

the ones closest to the nucleus

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7
Q

for electrons to move…?

A

… energy must be absorbed or released

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8
Q

as the shell number increases…?

A

… energy levels get closer

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9
Q

electron configuration

A

how electrons are arranged around the nucleus

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10
Q

Aufbau Principle

A
  • electrons enter orbitals of lowest energy

- orbitals within same energy level and sublevel all have equal energy

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11
Q

Pauli Exclusion Principle

A
  • an atomic orbital can contain at most 2 electrons
  • electrons in same orbital spin different directions
  • shown as arrows (up and down)- must have different spins
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12
Q

Hund’s Rule

A
  • every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied
  • all electrons in singly occupied orbitals have the same spin
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13
Q

filling order

A

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p

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14
Q

Noble Gas configuration

A

condensing electron configuration writing by using the closest previous noble gas

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15
Q

anion

A
  • negative ion

- add electrons to the next available orbital

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16
Q

cation

A
  • positive ion

- remove electrons from the last orbital

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17
Q

does the electron config change for ions

A

yes

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18
Q

what is the weird thing about ions in the d block

A

d group elements lose their 4s electrons before their 3d elements

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19
Q

isotopes

A
  • different forms of elements with a different mass because of added or removed neutrons
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20
Q

allotropes

A

different forms of the same element

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21
Q

relative atomic mass

A
  • the weighed average of the different isotopes of an element measured in a ratio
  • takes into account the percentage abundance of all isotopes that exist of an element
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22
Q

other names for relative atomic mass

A
  • atomic weight

- average atomic mass

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23
Q

Principal Energy Level (PEL)

A
  • designated by the quantum number n
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24
Q

relative atomic mass formula

A

look in book

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25
Q

bonding atomic radius

A
  • one half of the distance between covalently bonded nuclei

- the distance in picometres (pm) from the centre of the nuclei to the electrons in the outermost energy shell

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26
Q

electron shielding effect

A

the inner shell electrons repel the attraction of the valence electrons

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27
Q

atomic radius trend- group

A

as you go down a column, the atomic radius increases = stronger shielding effect

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28
Q

atomic radius trend- period

A

as you go across, L to R, atomic radius decreases = more electrostatic attraction (electrons put into same orbital)

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29
Q

smallest atomic radius’ in periodic table

A

noble gases

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30
Q

largest atomic radius’ in periodic table

A

alkali metal

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31
Q

ionic radius trend- metals

A
  • lose electrons = more attraction

- ionic radius < neutral atomic radius

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32
Q

ionic radius trend- non metals

A
  • gain electrons = less attraction

- ionic radius > neutral atomic radius

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33
Q

ionic size depends on?

A
  • nuclear charge
  • number of electrons
  • orbitals in which electrons reside
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34
Q

metals tend to form? (cations/anions)

A

cations

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35
Q

non metals tend to form? (cations/anions)

A

anions

36
Q

anions are ____ (larger/smaller) than their parent atoms

A

larger

37
Q

metallic characteristic trends

A
  • the ability to lose (or gain > non metallic) electrons
  • decrease from L to R across a period > caused by decrease in atom radius > allows valence electrons to ionize more readily
  • increase down a group > electron shielding causes atomic radius to increase > outer electrons ionize more readily than electrons in smaller atoms
38
Q

group 1 metals are highly reactive because?

A

they have one valence electron

39
Q

group 7 metals are highly reactive because?

A

they have seven valence electrons

40
Q

which ionic compounds are common

A

ones made from metals in group 1 and 7

41
Q

oxide trend

A
  • L to R - more basic > amphoteric > more acidic

- top to bottom - more acidic > amphoteric > more basic

42
Q

amphoteric

A

metal oxides that behave as both acidic oxides and basic oxides
(the more to the right the less these occur)

43
Q

ionization

A

losing or gaining an electron, electrons jump out of energy levels

44
Q

ionization energy

A
  • the amount of energy required to remove an electron from the ground state of a gaseous atom to form an ion
  • measured in kJ
  • first ionization energy = energy to move first electron, second ionization energy = energy to move second electron, etc.
45
Q

as ionization energy increases it becomes more ____ (easy/difficult) to remove an electron

A

difficult

46
Q

when all valence electrons have been removed…?

A

the ionization energy takes a quantum leap

47
Q

electronegativity

A

the tendency of the positive nucleus atom to electrostatically attract electrons

48
Q

a ____ (small/large) electronegativity can draw a neighbouring electron closer to itself in order to share the electron

A

large

49
Q

excited state

A
  • electrons gain energy and jump from their ground state to higher energy levels (further away)
  • the more energy absorbed, the further away it moves as it jumps
  • they become energetically unstable, and can fall back down > releases a certain amount of energy which would be seen as light at a particular colour or wavelength
50
Q

atomic emission

A
  • when electrons jump and cause different coloured lights as a result- the combined colour we see
51
Q

the greater the energy…?

A

the shorter the wavelength

52
Q

Concentration

A

Amount (mass)/volume

mg/L = mgL^-1

g/L

53
Q

electromagnetic spectrum

A

full range of frequencies of light

54
Q

excited state

A

when you heat an atom, the electrons gain energy and jump from ground state into any higher energy level, depending on how much energy is absorbed

55
Q

the more energy gained by electrons when heated….?

A

the further (or higher) they will jump, each jump involves a certain amount of energy

56
Q

what happens to electrons after they reach excited state

A

because they are now more energetic, when at unstable level they tend to fall back down to where they were before
> energy is released which can be seen as light of a particular colour or wavelength (UV or infrared)
- the electrons of the atoms in the flame can be promoted to higher orbitals for an instant by absorbing a set quantity of energy (a quantum)

57
Q

how is energy released from electrons returning to ground state

A

through light of a particular colour or wavelength (UV or infrared)

58
Q

when electrons jump, the greater the energy, the ____ (shorter/longer) the wavelength

A

shorter

  • big jumps = UV
  • medium = visible
  • small = infrared
59
Q

what can be seen as a result of all the electrons jumping?

A

a spectrum of coloured lines is produced > atomic emission > the colour you see is a combo of all the individual ones

60
Q

the amount of radiation from electrons jumping equals…?

A

the difference between the two energy levels

61
Q

why can there be more than one emission (electron jumps and falls)

A

because there can be more than one fall

62
Q

the amount of energy is specific to….? (electron jumping)

A

a particular electron transition in a particular element. As the quantity of energy is known and the quantity remaining can be measured, it is possible to calculate how many of these transitions took place, and thus get a signal that is proportional to the concentration of the element being measured

63
Q

types of transition series

A

Lyman, Balmer, Paschen

64
Q

Lyman Series

A
  • UV light

- transitions of electrons between the second (or greater) shell and the first shell

65
Q

Balmer Series

A
  • visible light
  • the transition of electrons in an atom from the third electron shell or greater, and the second shell
  • seen as coloured lines in an absorption and emission spectra
66
Q

Paschen Series

A
  • infrared light

- transition of electrons in an atom from the fourth shell or greater and the third shell

67
Q

AAS

A

atomic absorption spectroscopy

68
Q

AAS graphs

A

quantitative technique used to measure the concentration of metal atoms in a sample

  • in units given ppm or mg/L
  • sample is compared to a set of standards of known concentration of the metal being analysed
69
Q

ppm

A

parts per million

- same as mg/L

70
Q

calibration curves in AAS graphs

A

graph that plots concentration of standards on the x axis and absorbance on the y
- used to calculate unknown concentration

71
Q

as concentration of metal increases…?

A

the amount of light absorbed increases also

72
Q

concentration (metals)

A

amount or mass (mg)/volume (L)

73
Q

mgL^-1 can be also written as

A

mg/L

74
Q

calibration

A

find a value using the graph

75
Q

1ppm

A

1000 ppb

76
Q

1ppb

A

0.01ppm

77
Q

mass spectroscopy

A
  • technique used to measure the masses and relative intensities (abundances) of positive ions formed when a sample is bombarded with electrons and IV light
78
Q

what instrument is used to measure mass spectroscopy and what does it do

A

a mass spectrometer > used to determine relative isotopic masses and % abundance of elements

79
Q

mass spectrometer process

A
  1. using a flame or laser, rip the molecules in sample to shreds by electron bombardment. Generally this causes ionization - produces cations
  2. accelerate the ions toward a plate of negative charge > there should be a hole for the ions to pass through
  3. steer the ions toward another plate with even more negative charge. If the particles have lots of mass, they’ll travel further over the plate than particles with low mass because they have more momentum > particles then separated by mass
80
Q

the mass spectrum generated at the end of the mass spectrometer detects:

A
  • how many isotopes an element has, because each has a different mass-to-charge ratio
  • the relative mass of each isotope, indicated by the mass to charge ratio
  • how much of each isotope is in the sample, indicated by the percentage abundance
81
Q

how is the mass of each isotope measured (in mass spectroscopy)

A

comparing the mass of the isotope to the carbon-12 isotope (which is always a mass of 12). This is called relative isotopic mass.

82
Q

determining the relative isotopic mass and % abundance

A

on chart:

  • horizontal axis is the mass-to-charge ratio (m/z)
  • vertical axis is the % abundance of the sample
  • mass-to-charge ratio represents the RIM of an isotope
  • number of peaks = number of isotopes
83
Q

octet rule

A

all atoms want to have 8 valence electrons in their outer shell to have a similar form as noble gases, and this influences their bonding and they try to lose, gain, or share electrons to achieve this

84
Q

bonding

A

a bond is formed when valence electrons from two or more atoms/ions interact with each other, creating a strong electrostatic attraction

85
Q

types of bonding

A

ionic, metallic, covalent

86
Q

covalent bonding

A

between two nonmetallic elements of similar electronegativity, form molecules or covalent lattices
- formed by sharing electron pairs to satisfy the octet rule for a full shell of valence electrons

Year 11 Chemistry (1 decks)

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