Periodic Table Flashcards
(135 cards)
1
Q
What did Robert Boyle define an element as in 1660
A
A substance that cannot be broken into any simpler substances
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
3
Q
Elements are arranged in
A
Increasing order of atomic number
4
Q
Real name for mass number
A
Relative atomic mass
5
Q
Number of protons
A
Atomic number
6
Q
Vertical columns
A
Groups
7
Q
Horizontal rows
A
Periods
8
Q
How many man groups
A
8
9
Q
D-block made up of
A
Elements in groups II and III
10
Q
Short vertical columns in d-block
A
Sub-groups
11
Q
Group I
A
Alkali metals
12
Q
Group II
A
Alkaline Earth metals
13
Q
Group VII
A
Halogens
14
Q
Group 0
A
Noble gases
15
Q
Elements in 2 groups
A
Metal and non metals
16
Q
In general metals on…
A
Left of stairs
17
Q
In general non metals on…
A
Right of stairs
18
Q
Top step of stairs
A
B (5)
19
Q
Bottom step of stairs
A
At (85)
20
Q
Elemis bordering stairs…
A
Have similar properties to metals and non metals
21
Q
2 examples of elements bordering stairs
A
Si(14) and age (32)
22
Q
Most reactive metals
A
Group I alkalis
23
Q
Most reactive non-metals
A
Group VII halogens
24
Q
What did Dobreiner come up with
A
Dobreiners Law of Triads
25
What did Dobreiner note in his law of triads
Certain groups of 3 elements were related to their relative atomic mass
26
2 examples of triads
Lithium sodium potassium
| Sulfur selenium tellurium
27
Devi union of a triad
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
28
Ar
Relative atomic mass
29
What did Newland come up with
Newlands law of octaves
30
Newlands law of octaves
Each 8th element, starting from any given knee was similar in properties ti the first one
31
Where did newlands law of octaves work
For the first 16 elements
32
Why does newland slaw of octaves not work for the modern periodic table
Noble gases of group 0 are known now,
| It is now every 9th elements that is similar
33
How did newlands make an important contribution
He showed that the elements could be arranged in a table
34
What did Mendeleev come up with and how
Mendeleev’s periodic table
| He listed the known elements in order and put those with similar properties in vertical columns called groups
35
What gave rise to his periodic law
He noted that similar properties recurred periodically for every eighth element
36
What did Mendeleev do against his idea of increasing Ar and why
He put tellurium (Te) before Iodine (I) so that they would have similar properties to their groups
37
How was Mendeleev smart
He predicted the existence of many elements and left gaps for them and most were accurate eg. Germanium and Gallium
38
Mendeleev’s periodic law
When elements are arranged in order of Ar, their properties repeat at regular intervals or periodically
39
3 differences between Mendeleev’s and modern
Atomic number vs, Ar
No gaps in modern vs only 63 known back then
Number of blocks eg.d-block vs a rectangle
40
What provided mosely with an indirect method I’d measuring the number of protons in an atom
He noted that the frequencies of x-rays emitted by atoms of different elements varied with the quantity of positive charge (number of protons)
41
What did Mosley do with the periodic table
Put it in order of increasing atomic number and he showed that elements fell easily into their correct groups
42
At room temperature elements...
2 are liquid; mercury (Hg) and bromine (Br)
11 gases
Rest are solid
43
Nature of light
Consists of particles called photons which have energy but no mass and which travel in waves
44
Different colors are because of...
Different wavelengths, frequencies and energy contents
45
Read light
Long wavelength
Low frequency
Low energy
46
Violet light
Short wavelength
High frequency
High energy
47
How is a continuous spectrum formed
If white light is lasses through a prism as is dispersed a band of colores blend into each other
48
How is the emission spectrum of hydrogen formed
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
49
What does a line spectrum look like?
A few narrow band of light against a dark background
50
Emmisiom spectrum
The dispersed light from any source
51
Simplest emission line spectrum
From hydrogen gas
52
How are line spectrums unique?
Each element has its own emission spectrum which is different to that of any other element
53
Street lights
Sodium in a discharge tube
54
3 series in the emission spectrum of hydrogen
Lyman Balmer and paschen
55
Which series is ultra violet in
Lyman series
56
Which series is violet blue green and red in
The Balmer series (visible)
57
Which series is infra red in?
Paschen series
58
Electron in its lowest energy leve,
Ground state
59
Electron occupies...
Fixed energy levels
60
Moves from ground state up to...
Excited state
61
E=
E1-E2=hf
62
The fact that the line spectrum kr hydrogen consists of only a few lines of light of different energies shows...
The electrons of the hydrogen atom can only lose certain distinct amounts of energy and cannot lose a whole range I’d energies
63
The energy of the electron is said to be....
Quntisised
64
Why do we see all the lines at the same time in a hydrogen emisión spectrum
There are millions of hydrogen atoms in the discharge tube all doing millions of movements at the same time
65
Lyman series
A series of ultra violet lines caused by electrons falling back to n=1 level
66
Balmer series
A series of visible (violet, blue, green and red) lines caused by electrons falling back to n=2
67
Paschen
A series of infa red lines caused by electrons falling back to n=3
68
In Bohr’s theory of the atom what explained why electrons. Do not crash into the nucleus
The electron does not give out or take in energy unless it’s moving to another allowed energy level
69
The electron will move to a higher energy level if
It receives the exact energy equal to the difference between both energy levels
70
The electron will move to a lower energy level if
It loses an amount of energy exactly equal to tge difference between both energy levels
71
plancks constant
h
| 6.63 x10 -24 Js
72
f
Frequency of light emitted
73
Definition of energy level
A fixed or def8nite amount of energy that an electron is allergic to have in an atom
74
What causes elements to have unique emission line spectra
Each element has different numbers and different types I’d transitions
75
Electronic transitions
Movement from one energy level to another
76
How can you see an absorbtion spectrum
When whit slight is passed through a sample of it and the light is observed using a spectroscope
77
Why does it change when it goes through
It absorbs light of certain wavelengths which means that they don’t pass through
78
The absorbtion spectrum of an element is the...
Exact opposite of the emission spectrum
| It’s photographic negative
79
Absorbtion spectra are used in a laboratory technique called
Atomic absorption spectrometry
80
How do you measure the concentration of sodium using atomic absorbtion spectrometry
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
2 known uses for atomic absorbtion spectrometry
The analysis of water for lead and mercury
| In forensic science eg. Analyzing gun powder residue on clothes
82
Sublevels in n=1
1s
83
Sublevels in n=2
2s, 2p
84
Sublevels in n=3
3s, 3p, 3d
85
Sublevels in n=4
4s, 4p, 4d, 4f
86
Electrons in s
2
87
Electrons in p
6
88
Electrons in d
10
89
Electrons in f
14
90
Aufbau principle
Electrons must occupy the lowest energy levels available
91
Sub-energy level
A subdivision of a main energy level and consists of one or more orbitals of the same energy
92
What did Louis de Broglie state?
That all moving objects has a wave motion associated with that movement
93
Heisenberg’s uncertainty principle
It is not possible to determine at the same time the exact position and velocity of an electron
94
What did Schrodinger do? (Atomic orbitals)
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
Orbital
A region in space around a nucleus where there is high probability of finding an electron
96
The Pauli Exclusion Principle
Not more than 2 electrons can occupy an orbital and thrh cam only do this if they have opposite spin
97
Shape of s orbital
Sphere
98
Shape of px orbital
Dumbbell (horizontal)
99
Shape of py orbital
Dumbbell vertical
100
Shape of Pz orbital
Dumbbell
| On z axis (diagonal)
101
Hunds rule
When two or more orbitals of equal energy are available to electrons, the electrons will occupy these singly before occupying them in pairs
102
D-block metals
Elements which have their highest energy electrons in a d-sublevel
103
Transition metals
Meats which form at least one ion which has electronic configuration ending in an incomplete sublevel
104
Transition metal.., (3)
Have variable valiency (exist as different ions)
Exist as colored compounds
Act as catalysts
105
2 exception of aufbau principle
```
Copper
Chromium
Cu; Ends in 4s1, 3d10
Cr; ends in 4s1, 3d5
More stable ending in full or have sublevels
```
106
Why is there no yellow line in the line spectrum of hydrogen
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
Flame test result
| . Barium
Yellow/green
108
Flame test result
| . Calcium
Orange/red
109
Flame test result
| . Copper
Green/blue
110
Flame test result
| . Sodium
Yellow
111
Flame test result
| . Potassium
Lilac
112
Flame test result
| . Lithium
Crimson
113
Why are disposable wooden splints used in falm test
To prevent cross contamination
114
4 limitations of Bohr’s theory
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
Atomic radius
Half the distance between the nuclei of two atoms of the same element that are joined together by a single covalent bond
116
Atomic radius across a period
Decreases
117
Atomic radius down a group
Increases
118
Why does atomic radius decrease across a period
Increasing nuclear charge pulls outer electrons in closer to nucleus
119
Reasons for atomic radius increasing going down a group
```
Screening effect (more electrons)
Just bigger (more electrons)
(There is more nuclear charge but cancelled out by^^)
```
120
Why are group 1 more reactive going down the group (lose electrons when reacting) (2)
Atomic radius increases (further away from pull of nucleus)
| Increased screening
121
Why are group 7 elements live reactive doing down the group (gain electrons when reacting )
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
Nuclear charge
Number if protons (+) pulling in electrons (-) towards the nucleus
123
First ionisation energy
The energy needed to remove the most loosely bound electron from each atom in one mole or gaseous atoms in their ground state
124
First ionisation energy represent as (2)
X°(g)
Or
X+(g) + e-1
125
Second ionisation energy represented as (2)
X+ (g)
Or
X+2 (g) + e-1
126
Third ionisation energy represented as
X+2(g)
Or
X+3 (g) + e-1
127
Second ionisation energy
The amount of energy required to remove the most loosely bound electrons from each singly charged ion in one mole of gaseous ions
128
Ion ionisation energy going across a period...
Increases
129
Ionisation energy going doing a group
Decreases
130
Why does ionisation energy increase going a across a period (2)
Increased nuclear charge (stronger pull)
| Smaller atomic radius (closer to nucleus)
131
2 exceptions to general rule of ionisation energy increasing across a period
Be and N, very stable, full outer sublevels , high energy needs
Temporary decrease after these elements
132
Down a group decreasing ionisation energies
Larger atomic radius
| Extra screening
133
Successive ionisation energies of an element
The 1st, 2nd, 3rd, 4th etc. Ionisation energy of the same element (electrons being removed one by one)
134
When are successive ionisation energies increasing (3)
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
Except for spectroscopic experiments carried out by Bohr, what gives strong evidence that electrons have distinct energy levels?
The huge increases in some successive ionisation energies
