Chem Chap 1

Front 1

Proton, neutron electron (location, charge and mass)

Back 1

Proton: in nucleus, +1, 1AMU

Neutron: in nucleus, no charge, 1AMU

Electron: in orbitals around nucleus, -1, 1/1836AMU (negligible)

Front 2

What holds an atom together?

Back 2

The electrostatic attraction between the positively charged nucleus (protons have +ve charge)
and the negatively charged electrons.

Front 3

Deflection when passed between electrically charged plates:

Back 3

Like charges repel, unlike charges attract.

1. Electrons deflect away from negative plates towards positive.

2. Protons deflect away from positive towards negative.

3. Neutrons have no deflection.

(note: electrons undergo MORE deflection than protons cuz they are lighter in mass!!)

Front 4

Atom vs Ion

Back 4

Atom is neutral in charge (protons which have + are equal in number to electrons -)

Ions are charged atoms by gaining electrons becoming negatively charged

(anion) or losing electrons becoming positively charged (cation)

Front 5

Atomic Radius

Back 5

Measure of the size of an atom

Calculated as half the distance between the

nuclei of 2 covalently bonded atoms of the same element.

Eg: H2 is a H-H atom. If the distance between the nuclei of the 2 atoms is "d" then its atomic radius is "d/2"

Front 6

Atomic Radius trend across period

Back 6

Decrease across each period (atoms get smaller)
This is because a larger positive nuclear charge is being given but electrons are still in the same principal quantum shell. So there is a stronger positive pull by nucleus on the electrons= smaller atoms.

Front 7

Atomic Radius trend down group

Back 7

Increase down the group (atoms get bigger)

Number of shells increases and electrons repel each other and get farther away from the nucleus. Hence a bigger atom.

Front 8

Ionic radius

Back 8

measure of the size of an ion

Front 9

Trend of Ionic radii with increasing negative charge

Back 9

Increases since atom is gaining electrons to make an ion. Nuclear (positive charge) remained the same so the pull of the nucleus is less on the new electrons. Therefore an increase in ionic radius.

Front 10

Trend of ionic radii with increasing positive charge

Back 10

Decreases since here the atom is losing electrons to make an ion. The nuclear charge (positive) remained the same so the pull of the nucleus on the remaining electrons increased. Greater electrostatic force of attraction and hence a decrease in ionic radius.

Front 11

Isotopes definition

Back 11

Atoms of the same element that contain the same number of protons and electrons but a

different number of neutrons

Front 12

Symbol for different isotopes

Back 12

Chemical symbol of element followed by a dash and the mass number of that isotope.
Ex:
Carbon with mass number 12: C-12

Carbon with mass number 13: C-13

Carbon with mass number 14: C-14

All contain 6 protons (and electrons) but increase by 1 neutron each time.

Front 13

Physical and Chemical properties of isotopes

Back 13

Similar chemical properties due to the same number of outermost electrons (involved in reactions)

Different physical properties due to different neutron numbers which contribute to mass and density. (Density=mass/volume. if mass changes so does density)

Front 14

Electronic configuration meaning

Back 14

The arrangement of electrons in an atom

Front 15

Principal quantum shells. (how many electrons per shell)

Back 15

given the symbol (n) these are the energy levels around the nucleus.

The maximum number of electrons for each shell is:
n=1, 2e

n=2, 8e

n=3, 18e

n=4, 32e

(I know that you are used to saying 2,8,8,18 but this is the way it actually is, to find the number of electrons in a shell we use the formula 2n²)

Front 16

Subshells and orbitals

Back 16

s,p,d,f

S can carry 1 orbital= 2e

p can carry 3 orbitals=6e

d can carry 5 orbitals=10e

f can carry 7 orbitals=14e

Front 17

Ground state definition

Back 17

The most stable electronic configuration for an atom which has the lowest amount of energy.

Done by filling subshells with lowest energy first.

Pattern ends at n=3 but regular pattern is:
1s-> 2s-> 2p-> 3s-> 3p->4s->3d-> 4p-> 5s

(continuation not needed but is 5s-> 4d->5p->6s->4f->5d->6p-> 7s-> 5f-> 6d-> 7p-> 8s. This is why we say 2,8,8,18 do the math.

Remember, s s ps ps dps dps fdps fdps)

Front 18

Orbital energy

Back 18

Orbitals in the same subshell have the same

energy and are said to be degenerate. So p orbitals (known as px, py and pz) all have the same energy if they are in the same principal quantum shell.

Front 19

Quantum shell energy

Back 19

The energy of principal shells increase with increase in the number. n=4 has more energy than n=3 which has more energy than n=2 etc.

Front 20

Subshell energy

Back 20

Energy of subshells increases in the order s

Only exception is between 3d and 4s where 3d has more energy than 4s so the 4s orbital fills up first.

Front 21

S orbitals (shape and size)

Back 21

Spherical in shape and size increases with increase shell number. S orbital of n=3 is bigger than s orbital of n=1.

Front 22

P orbitals (shape and size)

Back 22

Dumbbell-shaped where every shell has a p orbital except n=1. The 3 orbitals in the p subshell occupy the x, y and z (vertical) axis and are at right angles to each other. The lobes (infinity symbol) get longer and larger as number of shells (n) increases.

Front 23

Electron configuration (numbers and their meanings)

Back 23

Electron configurations are written in the form of nsubshell power electrons held by subshell.

ex: 1s power 1 1s power 2 2s power 1 2s power 2 (maximum for s is 2) 2p power 1,2,3,4,5 or even 6.
EG: Mg with 12 electrons=

1s power 2 2s power 2 2p power 6 3s power 2

Remember that the power is the number of electrons. so 2+2+6+2=12

The number before the s or p only indicates the principal quantum shell we are in (n).

Front 24

Spin-pair repulsion

Back 24

Electrons in an orbital either spin clockwise or anticlockwise. Electrons that spin in the same direction will obviously repel each other causing a spin-pair repulsion. So to minimize this effect electrons try to fit in as many orbitals as they can.

As in if we have the electrons in a p subshell each electron will take a separate orbital rather than 2 electrons in 1 orbital and the last in the 2nd. The only time electrons are paired in an orbital is when there isn't any empty orbitals left. If this happens then electrons in an orbital will spin in OPPOSITE directions to minimize the repulsion. So if we have 5 electrons in a p subshell: 1 electron will be alone in an orbital and 4 will be paired up 2 by 2 in the 2 other orbitals spinning in OPPOSITE directions.

Check the diagram to know how to draw directions of spinning of an electron.

Front 25

Electron box notation

Back 25

A box represents the orbital. The boxes are arranged in order of increasing energy just like the subshells. Electrons are represented by opposite direction arrows to show that they are spinning oppositely.

Front 26

Electron box notation example

Back 26

Front 27

Free radical

Back 27

When there is 1 or more unpaired electron.

This unpaired electron is shown as a dot.
Example: Chlorine 17 is configured as 1s2 2s2 2p6 3s2 3p5. In the 3p5 subshell there are 4 electrons split in pairs into 2 orbitals while 1 electron remains unpaired in the 3rd orbital. This electron is known as the free radical electron and is represented by a dot Cl. (that's a dot for cl not the end of the sentence ). This free radical is formed by homolytic fission, when a covalent bond of 2 electrons is split evenly on 2 atoms. (Like in Cl2 being split to Cl alone with 17 electrons).

Front 28

Full Electron configuration

Back 28

Fully writing out the electrons in each subshell of an atom. As in Sodium would be:
1s2 2s2 2p6 3s1

Front 29

Shorthand electron configuration

Back 29

Using the nearest noble gas symbol BEFORE the element we are configuring then continuing whatever is left.

So sodium would be [Ne] 3s¹. Where Ne is neon which has 10 electrons. and the electron that's left is in the 3s orbital. So [Ne] 3s¹.

Front 30

Transition metals 4s and 3d subshells

Back 30

The transition metals fill the 4s subshell before the 3d subshell and lose electrons from the 4s first and not from the 3d subshell (the 4s subshell is lower in energy)

Front 31

s block d block d block f block elements

Back 31

s block elements have their valence electrons in an s orbital (first 2 groups)

p block elements have their valence electrons in a p orbital (Group 5,6,7,8)

d block elements have their valence electrons in a d orbital (transition elements. although they fill up the 4s orbital first and lose electrons from it first they still have their valence electrons in the d orbitals)

f block elements have their valence electrons in an f orbital (the lanthanides and actinides, the 2 periods underneath the periodic table)

Front 32

Exceptions to the configuration

Back 32

1. Chromium configured as
[Ar] 3d5 4s1 not [Ar] 3d4 4s2 (filled up only 1 electron in the 4s orbital instead of 2 before moving on to the 3d orbitals.)

2. Copper [Ar] 3d10 4s1 not [Ar] 3d9 4s2 (filled up only 1 electron in the 4s orbital instead of 2 before moving on to the 3d orbitals.)

This is because [Ar] 3d5 4s1 and [Ar] 3d10 4s1 configurations are energetically stable.

Front 33

Ionization Energy IE definition

Back 33

amount of energy required to remove one mole of electrons from one mole of gaseous atoms of an element to form one mole of gaseous ions.
In short the amount of energy required to remove one electron from one atom of a gaseous atom to form one gaseous ion of that atom.

Front 34

Ionization Energy standard conditions

Back 34

1. 298 K temperature

2. 1atm pressure

Front 35

Ionization Energy Unit

Back 35

kjmol-1

kj/mol

Front 36

Ionization Energies 1st 2nd 3rd

Back 36

The energies required to lose the first 2nd and 3rd moles of electrons respectfully from an atom to form 1st 2nd and 3rd mole of positively charged ions step by step and in order. (All in the gaseous state)

Front 37

Factors affecting 1st ionization energy

(4 points)

Back 37

Size of nuclear charge

Distance of outer electrons from nucleus

Shielding effect of inner electrons

Spin-Pair repulsion

Front 38

Ionization energy across a period

Back 38

Increases due to:

Nuclear charge increase making the atomic radii decrease and electrons get pulled closer and distance decreases. Shielding by inner electrons doesn't change as new electrons are being added to same shell. So more energy is needed to overcome the electrostatic attraction forces. So Ionization energy increases.

Front 39

Dips in trend (2 points)

Back 39

1. There is a slight decrease in ionization

energy between beryllium and boron as boron has a 2p subshell and hence is further away from the nucleus than the 2s subshell of beryllium.

2. There is a slight decrease between nitrogen and oxygen due to spin-pair repulsion in the 2px orbital (the first p orbital) of the oxygen atom. Since oxygen is 8 it has a configuration of 1s2 2s2 2px2 2py1 2pz1 so the 2 electrons in 2px2 have a spin-pair repulsion making it slightly easier for oxygen to lose that electron.

Front 40

Ionization energy from one period to the next

Back 40

Large decrease in ionization energy between last element in a period to first element in the next due to adding of new shell which causes:
increased distance (atomic radius)

increased shielding effect of inner electrons

These factors outweigh the +1 addition to the

nuclear charge.

Front 41

Ionization energy down a group

Back 41

Decreases, although the nuclear charge is increasing, the atomic radius is also getting bigger as distance increases due to adding of shells.

(every time you go down a group you add a shell)

Shielding effect of inner electrons also increases as shells are added.
These factors make it easier for the outer most electrons to be removed and so the ionization energy decreases.

Front 42

Ionization Energy down a group vs across a period

Back 42

Front 43

Successive ionization energy

Back 43

The loss of all electrons of an atom until the nucleus.
The successive ionization energy increases since:
a) every time you remove an electron the nuclear charge becomes greater than electron charge so its harder to lose electrons.

b) Atomic radius/distance decreases and shielding effect decreases as we move closer to the nucleus so more energy needed to overcome attractive forces.

Electrons aren't removed with periodic changes of energy however.

Front 44

Ionization energies, jumps and increases

Back 44

The first electron is always the easiest to

remove since its either the farthest away or is with another electron in the same orbital and has a spin-pair repulsion.

The second electron requires more energy if its in the same orbital as there is no spin-pair repulsion. But since they are in the same principal shell the increase in energy isn't that high.

There is a huge increase in energy required however as we move from one principal

quantum shell to another. This is because we get closer to the nucleus so:
Atomic radius/distance decreases and

shielding effect decreases.

Front 45

Summary of factors affecting ionization energy

Back 45