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Flashcards in Module 3.1 Deck (130)
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1

Give an outline of the periodic table

- Consists of rows (periods- horizontally across)
- Consists of groups
- Arranged by increasing atomic number

2

In the periodic table, elements are arranged by ...

Increase atomic number

3

What groups are fond in the s block?

Groups 1 and 2

4

What groups are found in the p block?

Groups 3-8

5

What are the names of the elements found in block d?

Transition metals

6

Define the term periodicity

Trends that occur in physical and chemical properties as we move across the periods of the periodict table

7

When talking about periodicity, list 5 different trends that we talk about

- Ionisation energy
- Melting Points/Boiling Points
- Reactivity
- Atomic Radius
- Electronegativity

8

Do all the elements within a group have similar reactions, why?

Elements in the same group have the same number of electrons on their outer shell. Therefore, they have similar chemical properties.

9

What was Döbereiner's theory about?

Triads

10

What was John Newlands' theory about?

Law of octaves

11

What was Mendeleev's theory about?

Gaps, elements arranged in order of increase atomic mass

12

What is the modern day theory of the periodic table?

Elements are arranged in order of increasing atomic number

13

Explain Döbereiner's theory behind his model of the periodic table

- Döbereiner grouped similar elements in TRIADS e.g. Li,Na and K.
- He ordered elements by atomic mass.
- The middle element in his triads had SIMILAR properties to the other TWO elements.

14

Explain John Newlands' theory behind his model of the periodic table

- Newlands arranged elements in order of mass. He noticed that every 8th element was SIMILAR.

15

Explain Mendeleev's theory behind his model of the periodic table

- He arranged all known elements by atomic mass
- He left GAPS, where no element fitted the repeating patterns
- Predicted the patterns of missing elements

16

Explain the theory behind the modern day periodic table ( Henry Moseley)

- Elements are arranged by increasing atomic number
- There are groups and periods

17

Define first ionisation energy

The amount of energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms.

18

Give the equation for the first ionisation energy of oxygen

O(g) → O^+ (g) + e^-

19

Give the equation for the first ionisation energy of magnesium

Mg(g) → Mg^+ (g) + e^-

20

Give the equation for the first ionisation energy of Sodium

Na(g) → Na^+ (g) + e^-

21

When talking about first ionisation energy what must you include?

- Charges
- Gaseous state
- The amount of energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms

22

The lower the ionisation energy, the easier/harder it is to form an ion?

Easier

23

What factors affect ionisation energy?

Nuclear charge
Atomic radius
Shielding

24

How does nuclear charge affect ionisation energy?

The higher the nuclear charge, the larger the electrostatic attraction between the nucleus and the outermost electrons. Therefore, more energy is required to remove 1 mole of electrons from 1 mole of gaseous atoms, so the as nuclear charge increase, so does IE.

Higher nuclear charge = Higher Ionisation energy

25

How does the atomic radius affect ionisation energy?

A increase in atomic radius means that the nuclear attraction between outermost electrons and the nucleus is smaller. Therefore, less energy is required to remove 1 mole of electrons from 1 mole of gaseous atoms, so as the atomic radius increases, ionisation energy falls.

Higher atomic radius = Lower Ionisation energy

26

How does the shielding affect ionisation energy?

As the number of electrons between the outermost electrons and the nucleus increases (As shielding increases), the outer electrons feel less attraction towards the nucleus. Therefore, less energy is required to remove 1 mole of electrons from 1 mole of gaseous atoms, so as shielding increases, ionisation energy falls.

More shielding = Lower ionisation energy.

27

Does Ionisation energy increase or decrease down the group, and why?

Down the group ionisation energy DECREASES. This is because elements down the group have extra electron shells compared to the ones above. The extra shells mean that the atomic radius is larger, so the outer electrons are further away (increasing shielding) from the nucleus, which greatly reduces their attraction to the nucleus.

28

Even though as you go down the group, the positive charge of the nucleus increases, why does ionisation energy still fall?

The effect caused by the positive charge is OVERRIDDEN by the effect of an increase in atomic radius and extra shielding

29

Does Ionisation energy increase or across the period, and why?

Across the period, ionisation energy INCREASES.

This is because across the period, the positive charge of the nucleus increases. This causes electrons to be pulled closer to the nucleus, making the atomic radius smaller. This means that the outer electrons are more strongly attracted to the nucleus. Therefore, more energy will be needed to ionise the element meaning across the period ionisation energy increases

30

Define the term second ionisation energy

The energy needed to remove an electron from each of one mole of 1+ ions in a gaseous state.