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Chemistry (Tim) > Periodic Table > Flashcards

Periodic Table Flashcards

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1
Q

What did Robert Boyle define an element as in 1660

A

A substance that cannot be broken into any simpler substances

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2
Q

What is one very important feature of the modern periodic table that makes the study of chemistry easier

A

Elements with similar properties are grouped together

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3
Q

Elements are arranged in

A

Increasing order of atomic number

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4
Q

Real name for mass number

A

Relative atomic mass

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5
Q

Number of protons

A

Atomic number

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6
Q

Vertical columns

A

Groups

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7
Q

Horizontal rows

A

Periods

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8
Q

How many man groups

A

8

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9
Q

D-block made up of

A

Elements in groups II and III

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10
Q

Short vertical columns in d-block

A

Sub-groups

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11
Q

Group I

A

Alkali metals

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12
Q

Group II

A

Alkaline Earth metals

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13
Q

Group VII

A

Halogens

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14
Q

Group 0

A

Noble gases

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15
Q

Elements in 2 groups

A

Metal and non metals

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16
Q

In general metals on…

A

Left of stairs

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17
Q

In general non metals on…

A

Right of stairs

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18
Q

Top step of stairs

A

B (5)

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19
Q

Bottom step of stairs

A

At (85)

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20
Q

Elemis bordering stairs…

A

Have similar properties to metals and non metals

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21
Q

2 examples of elements bordering stairs

A

Si(14) and age (32)

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22
Q

Most reactive metals

A

Group I alkalis

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23
Q

Most reactive non-metals

A

Group VII halogens

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24
Q

What did Dobreiner come up with

A

Dobreiners Law of Triads

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25
Q

What did Dobreiner note in his law of triads

A

Certain groups of 3 elements were related to their relative atomic mass

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26
Q

2 examples of triads

A

Lithium sodium potassium

Sulfur selenium tellurium

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27
Q

Devi union of a triad

A

A group of 3 elements with similar chemical properties where the relative atomic mass of the middle element is approximately the average of the other two

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28
Q

Ar

A

Relative atomic mass

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29
Q

What did Newland come up with

A

Newlands law of octaves

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30
Q

Newlands law of octaves

A

Each 8th element, starting from any given knee was similar in properties ti the first one

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31
Q

Where did newlands law of octaves work

A

For the first 16 elements

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32
Q

Why does newland slaw of octaves not work for the modern periodic table

A

Noble gases of group 0 are known now,

It is now every 9th elements that is similar

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33
Q

How did newlands make an important contribution

A

He showed that the elements could be arranged in a table

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34
Q

What did Mendeleev come up with and how

A

Mendeleev’s periodic table

He listed the known elements in order and put those with similar properties in vertical columns called groups

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35
Q

What gave rise to his periodic law

A

He noted that similar properties recurred periodically for every eighth element

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36
Q

What did Mendeleev do against his idea of increasing Ar and why

A

He put tellurium (Te) before Iodine (I) so that they would have similar properties to their groups

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37
Q

How was Mendeleev smart

A

He predicted the existence of many elements and left gaps for them and most were accurate eg. Germanium and Gallium

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38
Q

Mendeleev’s periodic law

A

When elements are arranged in order of Ar, their properties repeat at regular intervals or periodically

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39
Q

3 differences between Mendeleev’s and modern

A

Atomic number vs, Ar
No gaps in modern vs only 63 known back then
Number of blocks eg.d-block vs a rectangle

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40
Q

What provided mosely with an indirect method I’d measuring the number of protons in an atom

A

He noted that the frequencies of x-rays emitted by atoms of different elements varied with the quantity of positive charge (number of protons)

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41
Q

What did Mosley do with the periodic table

A

Put it in order of increasing atomic number and he showed that elements fell easily into their correct groups

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42
Q

At room temperature elements…

A

2 are liquid; mercury (Hg) and bromine (Br)
11 gases
Rest are solid

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43
Q

Nature of light

A

Consists of particles called photons which have energy but no mass and which travel in waves

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44
Q

Different colors are because of…

A

Different wavelengths, frequencies and energy contents

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45
Q

Read light

A

Long wavelength
Low frequency
Low energy

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46
Q

Violet light

A

Short wavelength
High frequency
High energy

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47
Q

How is a continuous spectrum formed

A

If white light is lasses through a prism as is dispersed a band of colores blend into each other

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48
Q

How is the emission spectrum of hydrogen formed

A

A sample of H2 gas that is través in a discharge tube is energized using electricity, it glows to give a faint light which is dispersed in a prism of a spectroscope

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49
Q

What does a line spectrum look like?

A

A few narrow band of light against a dark background

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50
Q

Emmisiom spectrum

A

The dispersed light from any source

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51
Q

Simplest emission line spectrum

A

From hydrogen gas

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52
Q

How are line spectrums unique?

A

Each element has its own emission spectrum which is different to that of any other element

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53
Q

Street lights

A

Sodium in a discharge tube

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54
Q

3 series in the emission spectrum of hydrogen

A

Lyman Balmer and paschen

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55
Q

Which series is ultra violet in

A

Lyman series

56
Q

Which series is violet blue green and red in

A

The Balmer series (visible)

57
Q

Which series is infra red in?

A

Paschen series

58
Q

Electron in its lowest energy leve,

A

Ground state

59
Q

Electron occupies…

A

Fixed energy levels

60
Q

Moves from ground state up to…

A

Excited state

61
Q

E=

A

E1-E2=hf

62
Q

The fact that the line spectrum kr hydrogen consists of only a few lines of light of different energies shows…

A

The electrons of the hydrogen atom can only lose certain distinct amounts of energy and cannot lose a whole range I’d energies

63
Q

The energy of the electron is said to be….

A

Quntisised

64
Q

Why do we see all the lines at the same time in a hydrogen emisión spectrum

A

There are millions of hydrogen atoms in the discharge tube all doing millions of movements at the same time

65
Q

Lyman series

A

A series of ultra violet lines caused by electrons falling back to n=1 level

66
Q

Balmer series

A

A series of visible (violet, blue, green and red) lines caused by electrons falling back to n=2

67
Q

Paschen

A

A series of infa red lines caused by electrons falling back to n=3

68
Q

In Bohr’s theory of the atom what explained why electrons. Do not crash into the nucleus

A

The electron does not give out or take in energy unless it’s moving to another allowed energy level

69
Q

The electron will move to a higher energy level if

A

It receives the exact energy equal to the difference between both energy levels

70
Q

The electron will move to a lower energy level if

A

It loses an amount of energy exactly equal to tge difference between both energy levels

71
Q

plancks constant

A

h

6.63 x10 -24 Js

72
Q

f

A

Frequency of light emitted

73
Q

Definition of energy level

A

A fixed or def8nite amount of energy that an electron is allergic to have in an atom

74
Q

What causes elements to have unique emission line spectra

A

Each element has different numbers and different types I’d transitions

75
Q

Electronic transitions

A

Movement from one energy level to another

76
Q

How can you see an absorbtion spectrum

A

When whit slight is passed through a sample of it and the light is observed using a spectroscope

77
Q

Why does it change when it goes through

A

It absorbs light of certain wavelengths which means that they don’t pass through

78
Q

The absorbtion spectrum of an element is the…

A

Exact opposite of the emission spectrum

It’s photographic negative

79
Q

Absorbtion spectra are used in a laboratory technique called

A

Atomic absorption spectrometry

80
Q

How do you measure the concentration of sodium using atomic absorbtion spectrometry

A

One would Energie’s a sample of pure sodium and allows the light to pass through the sample contaminated with sodium.
Only other sodium atoms will dully absorb the light while the others will reject it.

81
Q

2 known uses for atomic absorbtion spectrometry

A

The analysis of water for lead and mercury

In forensic science eg. Analyzing gun powder residue on clothes

82
Q

Sublevels in n=1

A

1s

83
Q

Sublevels in n=2

A

2s, 2p

84
Q

Sublevels in n=3

A

3s, 3p, 3d

85
Q

Sublevels in n=4

A

4s, 4p, 4d, 4f

86
Q

Electrons in s

A

2

87
Q

Electrons in p

A

6

88
Q

Electrons in d

A

10

89
Q

Electrons in f

A

14

90
Q

Aufbau principle

A

Electrons must occupy the lowest energy levels available

91
Q

Sub-energy level

A

A subdivision of a main energy level and consists of one or more orbitals of the same energy

92
Q

What did Louis de Broglie state?

A

That all moving objects has a wave motion associated with that movement

93
Q

Heisenberg’s uncertainty principle

A

It is not possible to determine at the same time the exact position and velocity of an electron

94
Q

What did Schrodinger do? (Atomic orbitals)

A

Used mathematical equations to predict where an electron might be found in spaces outside a nucleus
He plotted these points on 3-D polar diagrams

95
Q

Orbital

A

A region in space around a nucleus where there is high probability of finding an electron

96
Q

The Pauli Exclusion Principle

A

Not more than 2 electrons can occupy an orbital and thrh cam only do this if they have opposite spin

97
Q

Shape of s orbital

A

Sphere

98
Q

Shape of px orbital

A

Dumbbell (horizontal)

99
Q

Shape of py orbital

A

Dumbbell vertical

100
Q

Shape of Pz orbital

A

Dumbbell

On z axis (diagonal)

101
Q

Hunds rule

A

When two or more orbitals of equal energy are available to electrons, the electrons will occupy these singly before occupying them in pairs

102
Q

D-block metals

A

Elements which have their highest energy electrons in a d-sublevel

103
Q

Transition metals

A

Meats which form at least one ion which has electronic configuration ending in an incomplete sublevel

104
Q

Transition metal.., (3)

A

Have variable valiency (exist as different ions)
Exist as colored compounds
Act as catalysts

105
Q

2 exception of aufbau principle

A 
Copper
Chromium 
Cu; Ends in 4s1, 3d10
Cr; ends in 4s1, 3d5
More stable ending in full or have sublevels
106
Q

Why is there no yellow line in the line spectrum of hydrogen

A

Their is no corresponding energy loss by a hydrogen atom electron that releases light with a frequency or wavelength that would appear as yellow light

107
Q

Flame test result

. Barium

A

Yellow/green

108
Q

Flame test result

. Calcium

A

Orange/red

109
Q

Flame test result

. Copper

A

Green/blue

110
Q

Flame test result

. Sodium

A

Yellow

111
Q

Flame test result

. Potassium

A

Lilac

112
Q

Flame test result

. Lithium

A

Crimson

113
Q

Why are disposable wooden splints used in falm test

A

To prevent cross contamination

114
Q

4 limitations of Bohr’s theory

A

Only worked well for hydrogen
Could not explain splitting I’d certain emission line (sublevels)
Did not take into account the wave motion of the moving electron
At odds with Heisenberg’s uncertainty principle

115
Q

Atomic radius

A

Half the distance between the nuclei of two atoms of the same element that are joined together by a single covalent bond

116
Q

Atomic radius across a period

A

Decreases

117
Q

Atomic radius down a group

A

Increases

118
Q

Why does atomic radius decrease across a period

A

Increasing nuclear charge pulls outer electrons in closer to nucleus

119
Q

Reasons for atomic radius increasing going down a group

A 
Screening effect (more electrons)
Just bigger (more electrons)
(There is more nuclear charge but cancelled out by^^)
120
Q

Why are group 1 more reactive going down the group (lose electrons when reacting) (2)

A

Atomic radius increases (further away from pull of nucleus)

Increased screening

121
Q

Why are group 7 elements live reactive doing down the group (gain electrons when reacting )

A

Atomic radius increases, new electrons can’t get close to nucleus to be held tightly
Extra screening makes it more difficult to gain electrons

122
Q

Nuclear charge

A

Number if protons (+) pulling in electrons (-) towards the nucleus

123
Q

First ionisation energy

A

The energy needed to remove the most loosely bound electron from each atom in one mole or gaseous atoms in their ground state

124
Q

First ionisation energy represent as (2)

A

X°(g)
Or
X+(g) + e-1

125
Q

Second ionisation energy represented as (2)

A

X+ (g)
Or
X+2 (g) + e-1

126
Q

Third ionisation energy represented as

A

X+2(g)
Or
X+3 (g) + e-1

127
Q

Second ionisation energy

A

The amount of energy required to remove the most loosely bound electrons from each singly charged ion in one mole of gaseous ions

128
Q

Ion ionisation energy going across a period…

A

Increases

129
Q

Ionisation energy going doing a group

A

Decreases

130
Q

Why does ionisation energy increase going a across a period (2)

A

Increased nuclear charge (stronger pull)

Smaller atomic radius (closer to nucleus)

131
Q

2 exceptions to general rule of ionisation energy increasing across a period

A

Be and N, very stable, full outer sublevels , high energy needs
Temporary decrease after these elements

132
Q

Down a group decreasing ionisation energies

A

Larger atomic radius

Extra screening

133
Q

Successive ionisation energies of an element

A

The 1st, 2nd, 3rd, 4th etc. Ionisation energy of the same element (electrons being removed one by one)

134
Q

When are successive ionisation energies increasing (3)

A

Always increasing
Very large increase from one main energy level to another eg. n=2 and n =3
Higher when taking from a full sublevel eg. 2p6

135
Q

Except for spectroscopic experiments carried out by Bohr, what gives strong evidence that electrons have distinct energy levels?

A

The huge increases in some successive ionisation energies

Chemistry (Tim) (8 decks)

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