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Atomic Structure Flashcards

1
Q

What does the principal quantum no indicate

A

The shell occupied by the electrons

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2
Q

What is a shell

A

A group of orbitals with the same principal quantum number

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3
Q

How many electrons can the first second third and fourth shell hold

A

2,8,18,32

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4
Q

What is an orbital and what are the diff types

A

Region around the nucleus that can hold upto two electrons with opposite spins
S
P
D
F

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5
Q

What’s the shape of the s and p orbitals and how many diff orbitals are there of each

A

S-spherical
P-Dumbell shaped
S=1,p=3,d=5,f=7

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6
Q

What are the rules by which electrons are arranged in a shell?

A

-Electrons are added one at a time
-Lowest available energy level filled first
-Each energy level must b filled before the next one
-Each orbital filled singly before pairing
-4s filled before 3d

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7
Q

Why does 4s original fill before 3d?

A

4s orbital has a lower energy level than 3d before it is filled

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8
Q

What’s the electronic configuration for iron?
(26 electrons)

A

1sq,2sq,2p6,3s2,3p6,4s2,

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9
Q

What’s the electronic configurations for copper and chromium and state why we do this? (Cu-29)
(Cr-24)

A

Copper -1s2,2s2,2s6,3s2,3p6,4s1,3d10
Cu-1s2,2s2,2p6,3s2,3p6,4s1,3d5
The 3d subshell is more stable when it’s half full or completely full

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10
Q

What’s the shorthand electronic configuration of sodium? (11 electrons)

A

[Ne]3s1

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11
Q

Which electrons are lost when atoms become a positive ion

A

Electrons in the highest energy levels

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12
Q

What’s the electronic configuration of a sulfur ion?

A

1s2,2s2,2p6,3s2,3p6

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13
Q

If the iron atom has the electronic conguration of 1s2,2s2,2p6,3s2,3p6,3d6,4s2
Then what is the electronic configuration for the iron 2+ ion?

A

1s2,2s2,2p6,3s2,3p6,3d6,

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14
Q

If the iron atom has the electronic conguration of 1s2,2s2,2p6,3s2,3p6,3d6,4s2
Then what is the electronic configuration for the iron 2+ ion?

A

1s2,2s2,2p6,3s2,3p6,3d6,

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15
Q

What’s the definition of the first ionisation energy?

A

The energy needed to remove one mole of electrons from one mole of atoms in their gaseous state to form one mole of 1+ions (also in gaseous state).

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16
Q

What’s the first ionisation energy for magnesium?

A

Mg—-mg+ + e-

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17
Q

What’s the definition for the second ionisation energy

A

The energy needed to remove one mole of electrons from one mole of 1+ ions in their gaseous state to form one mole of 2+ ions(also in their gaseous state)

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18
Q

What’s the fourth ionisation energy equation for mg

A

Mg3+ ——mg4+ + e-

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19
Q

What are successive ionisation energies?

A

When we can continue to remove electrons and measure the ionisation energies

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20
Q

Describe the attraction between nucleus and electrons when it comes to ionisation energy

A

The electrons in an atom are attracted to the positive protons in the nucleus.The greater the attraction between the outer electrons and the nucleus the greater the ionisation energy.

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21
Q

What are the 3 factors that affect ionisation energy?

A

The atomic radius/distance between nucleus and outermost electrons. As the AR increases the force of attraction between the positive nucleus and outer electrons decreases.
The greater the no of protons the greater the force of attraction between the outer electrons and the nucleus.
Shielding-electrons in the outer shell are repelled by electrons in inner shells and this shielding effect reduces the attraction between the outer electrons and the nucleus.

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22
Q

|. *
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|. *
|. *
|. *
| *
|. *
|*
|___________________________________
1st 2nd 3rd 4th 5th 6th 7th 8th
Explain the gradual increase in this ionisation graph?
And why the 7th is much higher

A

Each time we remove an outer electron the remaining electrons in the outer shell are pulled slightly closer to the nucleus.this means there’s a greater attraction to the outer electrons and nucleus and this causes ionisation energy to increase.

After we removed the 6 electrons in the outer shell the 7nth was in the inner shell closer to the nucleus and these have much less shielding,this means they have a greater attraction to the nucleus and why the ionisation energy is much higher

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23
Q

|. *
| *
|
|
|
|.
|.
|. *
| *
|. *
|*
|___________________________
1st 2nd 3rd 4th 5th 6th
Identify the element?

A

It’s in group 4

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24
Q

Why does the first ionisation energy decrease as we go down a group?

A

The atomic radius increases, this means that the outer electron shell is further away from the nucleus.
Going down the no of internal energy levels increases so more shielding between electrons and nucleus
-> both of these causes the attraction between outer electrons and nucleus to decrease causing first ionisation energy to fall.

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25
|. | | | | |. |. |. |. *be *c *N |*Li *B *O |___________________________________ 1st 2nd 3rd 4th 5th 6th 7th 8th Explain the increases and the decreases in this graph?
As we move across the nuclear charge increases as the no of protons increases. This increases the attraction between the nucleus and the electrons.So atomic radius decreases across a period.increased nuclear charge and decreased atomic radius increase the attraction between outer electrons and nucleus— causing first ionisation energy to increase. For the others they have the same electronic configuration we’re removing the electron from. For boron we’re removing it from the 2p shell as this has a higher energy than the 2s subshell and further away from nucleus so this takes less energy to remove the outer electrons of boron causing ionisation energy to fall. For oxygen there’s a pair of spin electrons in the 2p subshell and as these electrons repel each other it takes less energy to remove these electrons than if they were in seperate orbitals.
26
What is an ionic bond Why do many atoms react to form an ionic bond And between what type of atoms does it occur
electrostatic force of attraction between posoitve and negative ions To achieve the electronic configuration of a noble gas Metal to non metal
27
What do we form when we react metal and non metal
An iomic compound giant ionic lattice as each ion attracts oppositely charged ions in all directions This attraction is the electrostatic force
28
Properties of ionic bond
High melting point and boiling point Dissolves in polar solvents like water Conducts entrichtet when liquid or aq
29
Why do ionic bonds have a high mp and bp How can melting point be higher for one atom
great deal of energy needed to overcome the strong electrostatic forces of attraction this can only be done by high temp If they’ve got higher ionic charges w stronger elec forces of attraction between ions or smaller ions are closer tg so greater force of attraction.
30
Why can ionic compounds dissolve in polar solvents
Polar water molecules can break down the ionic lattice and then the positive end of water molecule will attract to the negative ion and the negative end of a water molecule will attract to the positive ion. This overcomes electrostatic forces between ions so they dissolve in polar solvents If a compound has higher charges the ionic attraction too strong so then the water molecules can’t overcome the electrostatic forces of attraction and break down the lattice structure. Solubility of compound decreases
31
Why don’t ionic compounds conduct electricity
When solid ions are fixed in giant ionic lattice so there are no moving charge carriers When liquid the lattice is broken down and ions are free to move and carry charge
32
What is covalent bonding what does it occur between And why does it react with a covalent bond
The strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms. Non metallic elements compounds of the non metallic elements and polyatomic ions To achieve the same electron structure as the nearest noble gas in periodic table
33
What happens between boron when it comes to covalent bonding
Boron has only three electrons in outer shell to be paired. When these are paired it forms boron trifluroide but only had 6 atoms in outer shell so it doesn’t achieve same elec config as noble gas
34
What happens to phosphorous sulfur and chlorine
Phosphorus has 5 electrons in outer shell and when these are paired they form phosphorus trichloride to get same elec config as noble gas but this leaves two electrons without a shared pair so this is the lone pair of electrons. When all the pairs are shared to make phosphorus pentaclhoirde it now has 10 electrons in outer shell but because phosphorus in group 3 so it’s in the D subshell which the max it can have is 18.
35
What’s it called when you use the d subshell in that way to create more shared pairs
Expansion of the octet
36
When do multiple covalent bonds exist?
When two atoms share more than one pair of electrons
37
What happens in double covalent bonds
The electrostatic force of attraction is between two shared pairs of electrons and the nuclei of the bonded atoms All atoms have a full outer shell of 8 electrons and same electronic configuration to the nearest noble gas.
38
What are triple covalent bonds
The electrostatic attraction is between three shared pairs of electrons and the nuclei of the bonding atoms. So all atoms have an elec config as a noble gas.
39
What a dative covalent bond
Covalent bond in which the shared pair of electrons has been supplied by one of the bonding atoms only. This would originally be a lone pair of electrons on one of the bonded atoms
40
How does the formation of nh4 ion work and how do your display it
Nh3 donates a a lone pair of electrons to H as it has no electrons It uses its lone pair of electrons to form a covalent bond to the hydrogen ion forming nh4 H | | H—|——>H | | H To show nitrogen atom provides both electrons to the covalent bond
41
What’s average bond enthalpy
It serves as a measurement of covalent bond strength. The larger the value of the average bond enthalpy the stronger the covalent bond.
42
What is electron repulsion theory
The shape of a molecule is determined by the electron pairs surrounding the central atom. Based only on outer shells And based on the fact that pairs of electrons repel all other pairs of electrons and electrons move as far boat to. minimise the repulsion
43
What is a pair of electrons
A covalent bond
44
What structure and angle does beryllium chloride and carbon dioxide have
Beryllium chloride is covalently bonded to other chlorine atoms with two electron pairs. These repel each as much as possible in a linear structure and the angle between the bonds is 180c• In co2 the central atom is carbon which is covalently bonded to two oxygen atoms with double bonds Double bonds treated same as single So these repel each other as far as possible in a linear structure with an angle of 180c•
45
What structure and angle does boron trifluoride have
It’s central atom is boron which is covalently bonded to three atoms of fluorine the electron pairs in these three repel and move apart as far as possible in a triagonal shape and it’s planar as it’s viewed flat So it’s triagonal planar with an angle of 120c•
46
When would it usually be triagonal planar
When you’ve got a central atom with three pairs of bonding atoms around it and central atom has no line pairs around it
47
What structure and angle is it when the central atom has 4 pairs of electrons around it
Tetrahedal with 109.5c• Eg ammonia or methane
48
What structure and angle is it when the central atom has 5 pairs of electrons around it
To minimise repulsion two of the bonding pairs move to opposite sides of the molecule The other three bonding pairs take up a central position lying on the same plane The bonds pouting up and down is at 90c• and the one lying on central plane is 120c• This is trioganal bipyriamidal
49
What do solid lines in shapes of molecules represent A solid wedge? Dotted wedge
A bond in the plane of the paper Comes out of the plane of the paper Goes into plane of paper
50
What’s the effect of lone pairs
They repel more strongly than bonded pairs and this extra repulsion decreases other angles by 2.5c•
51
In ammonia the nitrogen atom has a lone pair what’s it’s structure and plane
It has 4 pairs of electrons so it’s a tetrahedral structure with 109.5c• but a lone pair repels more strongly than a bonding pair so this decreases the angle by 2.5 making it 107c• i and this is a pyramidal shape as it looks like a pyramid
52
Whattt happens when ammonia forms an ammonium ion
When the ammonia bonds to a hydrogen ion the lone pair forms a dative covalent bond . This has the same level of repulsion as a regular covalent bond soo bond angle in ammonium ion returns back to tetrahedral angle 109.5
53
What happens with Water which has Single covalent bonds to hydrogen but 2 lone pairs on the oxygen
As oxygen has 2 lone pairs the total number of bonding areas is 4 so it’s a tetrahedral with an angle of 109.5. As lone pairs repel stronger they decrease by 2.5c• and as there’s two lone pairs it’s by 5c• making the angle 104.5. This is a non linear molecule
54
If it had 6 pairs of bonded electrons what structure and angle is it
Octahedral 90c•
55
What is electronegativitu
The ability of an atom to attract the pair of electrons in a covalent bond
56
In a h2 molecule where do the electrons lie and what’s the significance of it
Right in the middle of the nuclei of both because they both have the same electronegativity so pair of electrons is equally attracted to the two Hydrogen nuclei
57
In a hydrogen chloride molecule where does the electron pair lie and what’s the significance of it
Closer to the chlorine nucleus because it’s more electronegative
58
What’s a property of simple molecular substances
Low bp- when we heat sms the molecules move faster and the intermolecular forces are easily broken down as little energy is needed
59
What are the different types of intermolecular forces
•Induced dipole dipole interactions(London) •Permanent dipole dipole interactions •Hydrogen bonds
60
How does an atom have an induced dipole
If the electrons in an atom are randomly moving and for a fraction of a second favour one side of an atom (positive or negative ) it then has a dipole and it’s instantaneous because of random electron movement Because it’s moved to one side this causes a dipole on another atom as the electrons repel each other making it induced
61
What’s another word for induced dipole dipole interaction that the atoms experience
An attraction called a London force- these are weak and easily broken ( it’s an IM force so weaker than covalent) Every single atom experience London forces as they’re made due to random electron arrangement The strength of London forces is due to the number of electrons
62
How do permanent dipole dipole interactions work
When one atom has greater electronegativity than another the electron pairs are more attracted to that making it slightly more negative. When two molecules of this are put together they experience a perm ant dipole dipole interaction
63
How can an atom have a permanent dipole
When there’s imbalance of charge distribution so if the charges are symmetrical it doesn’t have a permanent dipole
64
Trichloromethane and tetrachloromethane has bp of 61 and 77 Explain why tetra has a higher one if it doesn’t have a permanent dipole
Because permanent dipole aren’t the pnly forces actinf as London forces act as well and London forces increase based on the number of electrons and tetra has higher so it has a higher bp
65
What is the strongest type of IM
Hydrogen bonds
66
What are the conditions for hydrogen bonding to occur
A hydrogen atom attached to a strongly electronegative element The electronegative element must have a lone pair of electrons
67
What elements could this be
Fluorine oxygen and nitrogen
68
Water has a boiling point of 100c• and hydrogen Sulfide has a boiling point of -60c• explain why
In water the Hydrogens are bonded to oxygen which is the second most electronegative element. This means water molecules undergo hydrogen bonding. As it’s the strong IM force it takes a lot of energy to overcome so it had a high bp. Hydrogen sulfide the Hydrogens Are linked to sulfur which is as electronegative than hydrogen so it can’t form a hydrogen bond.
69
Why is ice less dense than water
Due to hydrogen bonding, so when it’s liquid the water molecules are moving randomly moving further apart and closed tg always breaking and reforming. At low temps water freezes (L-S) water molecules are held further apart in ice than they are in liquid water because each water molecule forms four hydrogen bonds to other water molecules forming a lattice shape. This makes ice less dense than water hence why ice floats.
70
Why is this beneficial
Because ice insulates the water below preventing animals from completely freezing
71
Draw a diagram of hydrogen bonding
72
Describe bonding in simple molecular substances
Atoms within the same molecule are held by strong covalent bonds and different molecules are held by weak intermolecular forces.
73
Why do simple molecular structures have low bp and mp
Small amount of energy needed to over come the weak IM forces
74
Can simple molecular substances conduct electricity and why
No because there are no free charged particles to move around
75
When is a molecule non polar
When both atoms have the same electronegativity
76
When a simple molecular substances freezes below it’s mp what happens
It forms a simple molecular lattice( all halogens do this btw)
77
What do non polar substances dissolve in and what do they not dissolve in
Non polar solvents Polar solvents eg water because it’s hydrogen bonded to other molecules so non polar solvents become insoluble
78
Where does polar substances dissolve in
polar solvents ( highly soluble)
79
What’s two anomalous properties of water
Ice being denser than water Ice having a high np
80
When a molecule ha s a bond angle why would this be
Due to the electron pairs repelling each other as far as possible
81
If a molecule has two bonded atoms and two lone pairs what is it
Non linear 104.5
82
If a molecule has 2 bond pairs and no line pairs what is it
Linear -180
83
How does a polar bond occur
When one atom has a higher electronegativity than the other so the shared pair of electrons are shared unequally making one more negative
84
How does a

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