CfE Higher Chemistry Unit 1: Chemical Changes and Structure

Trends in the Periodic Table and Bonding

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Learning Outcomes :

Arrangement of Elements in the Periodic Table

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•I can predict the properties of an element using its position in the Periodic Table•I can identify groups and periods in the Periodic Table.•I can explain why certain elements have similar properties•I can identify the alkali metals, halogens, noble gases and transition elements in the Periodic Table.

Lesson Starter: Data Book Task

. Use your iPad to join the Higher Chemistry class on iTunesU.Enrolment key for the iTunesU course is:

FKH-NPF-BXL

. Download a copy of the Higher Chemistry Data book either from the SQA website or the iTunesU resource folder.

. Using your data book and using your knowledge of chemistry comment on the use of titanium metal in the aerospace industry.

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The Periodic Table• On March 6th 1869 Dmitri Mendeleev, a Russian chemist,

published his Periodic Table of the Elements. He arranged the known elements in order of increasing atomic masses. This was eventually changed to atomic number.

• Elements with similar chemical properties were arranged in groups. He left gaps for elements yet to be discovered

• In the years that followed his ideas were modified and new elements were discovered until we arrived at the modern Periodic Table.

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Trends in the Periodic TableThe elements in the periodic table have different properties

The table is set up in such a way that these properties vary periodically

across a period,

or down a group

The properties are both physical and chemical

The chemical properties of an element stem from its physical properties:

Density, melting points and boiling points, atomic size, ionisation enthalpy, attraction for bonding electrons

Notes:

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After discussion with your group, make sure you can identify where all the following groupings are in the Periodic Table and what their properties are. Mark these on the periodic table handout and stick it into your notebook. Metals Non-metals Alkali metals

Transition metals Halogens Noble gasesThe diatomic elements The radioactive elements

Notes:

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Metals Non-metals

Alkali metals

Transition metals

Halogens

Noble gasesThe diatomic

elements

The melting and boiling points of elements give an indication of the forces that hold the atoms or molecules together

The higher the melting and boiling point the stronger the forces

The trend is similar for both melting and boiling so we’ll just look at melting

Melting and Boiling points

Notes:There are variations in the physical properties of the elements

across a period and down a group.(i) Density• Copy and complete the table below for the first twenty

elements, using the information in the data booklet.

• What do all the elements in Group A have in common?• What do all the elements in Group B have in common?• How does the density of the elements change across a period?• How does the density of the elements change down a group?

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Physical properties of the elements

Variation of density (g cm-3) with atomic number

period 2 (Li - Ne) maximum at boron (B) - group3period 3 (Na - Ar) maximum at Aluminium (Al)- group 3

Variation of density (gcm-3) with atomic numberIn general in any period of the table, density first increases from group 1 to a maximum in the centre of the period, and then decreases again towards group 0

2nd 3rd

4th

5th

Variation of density (g cm-3) with atomic numberAdapted from New Higher Chemistry E Allan J Harris

down a group gives an overall increase in density

Notes:(ii) Melting and boiling points• Using the information in the data booklet, how do the meltingand boiling points of the elements change across a period?• What is the trend in melting and boiling points down Group 1?• What is the trend in melting and boiling points down Group 7 and

Group 0?

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• The melting point starts off low, gradually increases to a peak (at group 4) then gradually decreases to a very low value (at group 0 or 8)

• To explain this trend we must think about the strength of the forces between the molecules

• In group 1 the atoms are held together by metallic bonds• In group 4 the atoms are held together by many very strong covalent bonds

(covalent network)• In group 8 the atoms are held together by very weak bonds (monatomic

gases)• We will look at the different types of bonding later in the unit

Learning Outcomes :

Bonding and Structure of the first twenty elements

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•I can explain how a covalent bond is formed.•I can describe the behaviour of outer electrons in metallic bonding.•I can explain the difference between covalent network and covalent molecular.•I can give examples of metallic, covalent molecular, covalent network and monatomic elements.

Lesson Starter: N5 2014 PP Question

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Bonding in the first 20 elements

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Demo 1.6

1. Metallic Bondinge.g. Li, Na, K, Be, Mg, Ca, Al

• Strong electrostatic forces exist between the positive nuclei and the outer shell electrons.

• These electrostatic attractions are known as metallic bonds.

Positive nucleus (core)Delocalised electron

+ + + +

+ + + +

• The outer shell in metals is not full and so metal electrons can move between these partially filled outer shells.

• This creates what is sometimes called a ‘sea’ or ‘cloud’ of delocalised electrons.

Physical properties of metals1. Metals are malleable and ductile

2. Conduction of electricity

Metal atoms can ‘slip’ past each other because the metallic bond is not fixed and it acts in all directions.

The ‘sea’ of delocalised electrons can move and carry the charge+

applied force

M.p.’s are relatively low compared to the B.P’s. This is because in a molten metal the metallic bonding is still present.B.p.’s are much higher as you need to break the metallic bonds throughout the metal lattice.

3. Change of state

Metal b.p.’s are dependant on(i) How many electrons are in the outer shell(ii) How many electron shells there are.

2. Covalent Networks

Giant lattice of covalently bonded atoms.e.g. B, C, SiGiant molecules held together by covalent bonds, resulting in high mpt and bpt.

Boron: M.pt.= 2573K= 2300oCSilicon: M.pt.= 1683K= 1410oC

Model 1.8

Two Forms of Carbon

1. Graphite (covalent network)

Weak Van der Waals forces between layers

Strong covalent bonds between atoms

• 3 electrons from each C atom used in bonding.

• 1 electron from each C is delocalised so graphite conducts electricity.

• Stacked hexagonal layers of C atoms with only weak Van der Waals between layers so layers can slip and slide- graphite is soft and can be used as a lubricant.

2. Diamond (covalent network)• Hardest natural substance à very strong covalent

bonds.• Electrical insulator (no delocalised electrons)-all outer

electrons are localised in covalent bonds.• Tetrahedrally bonded carbon atoms

We will discuss the 3rd form of

carbon later

3. Discrete Covalent Molecules• These molecules have known numbers of

atoms (discrete molecules).e.g. Oxygen M.pt.= 55K= -218oC

Sulphur M.pt.= 386K= 113oC

• Low Mpts indicate that weak Van der Waals forces are present.

a) Diatomic molecules

H – H O = ON ≡ NF – FCl – Cl

Remember HON 7!

à All gases due to weak Van der Waals forces.

Diatomic

b) P and S

Covalent solids held together by Van der Waals which are stronger due to higher molecular masses.

c) Carbonà Buckminster Fullerene• Very large, C60.

4. MonatomicGroup 8 elements e.g. He, Ne, Ar

Non-bonded atoms. Only weak Van der Waals forces between atoms.

Noble gases have full outer shells, they do not need to combine with other atoms.

Noble gases

b.p / oC

B.p.’s increase as the size of the atom increasesThis happens because the Van der Waals forces increase

-280-260-220-200

-180-160-140-120-100

HeliumNeonArgonKryptonXeon

Practice Question PP Question 2004

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Bond Strengths

Bond Type Strength (kJ mol –1)Metallic 80 to 600

Ionic 100 to 500Covalent 100 to 500

Van der Waals forces

1 to 40

Learning Outcomes :

Bonding in compounds05/09/2017

•

Bonding and melting point

substance ionic or covalent molecular did it melt?

copper sulphatesalol

potassium nitratesodium chloride

paraffin wax

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Expt 1.9The properties of a compound depend on the type of bonding present.In the following experiment you will investigate melting points of ionic and covalent molecular solids. Place the test tubes provided in a beaker of boiling water for a few minutes

Find out the actual melting point of these compounds.

Ionic compounds have ……………. melting points. • Explain this in terms of arrangement and movement of

particles as well as attraction between particles.

Covalent molecular solids have …………………… melting points.

• Explain this in terms of attraction between and movement of particles.

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Covalent network compoundsSilicon dioxide has the formula SiO2. Silicon carbide has the

formula SiC.• What type of bonding would you expect to exist in these compounds?• Would you expect these compounds to have high or low melting points?• Find out the actual melting points.• How do they fit with your prediction?• Examine models of these compounds.• Explain melting point and boiling point in terms of bonding and movement

of particles• Consider the valances of C and Si and use this information to work out the

exact structure of silicon carbide.• Silicon carbide (SiC) is widely used as an abrasive as it is an extremely hard

material. Explain this is terms of its structure.

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Success Criteria:

Next Lesson:

ü I can explain how a covalent bond is formed.ü I can describe the behaviour of outer electrons in

metallic bonding.ü I can explain the difference between covalent

network and covalent molecular.ü I can give examples of metallic, covalent

molecular, covalent network and monatomic elements.

Patterns in the periodic table: Covalent Radius

Learning Outcomes :

Periodic Trends in Ionisation Energy and Covalent radius

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•I can use covalent radius to describe the changes in the size of atoms across a period and down a group.•I can explain the change in covalent radius in terms of changes in the number of occupied shells or the nuclear charge.•I can state what is meant by first, second and third ionisation energies.•I can write state equations to represent first, second and third ionisation energies.•I can use atomic size and screening effect to explain the change in ionisation energies down a group.•I can use atomic size and nuclear charge to explain the change in ionisation energies across a period.

Covalent Radii of ElementsThe size of an atom is measured by it’s covalent radius, the distance between the nucleus and it’s outer electrons.

Values for covalent radii can be found in the data book

nucleus

energy levels

covalent radius

Looking down a group

The single electron in the outermost energy level is much further from the nucleus in caesium.

CsLi

-

-

This causes the caesium atom to have a much larger covalent radius.

The caesium atom also has many more electrons between the single outer electron and the nucleus.

This screening effect counteracts the attraction from the greater nuclear charge.

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Looking across a period

Across a period we can see the covalent radius decreasing.

So, from lithium to fluorine:

3+

-

9+

-

--

--

-

-

Lithium Atom Fluorine Atom

As we move left to right we are adding a proton to the nucleus and an electron to the outermost energy level.

Looking across a period

The lithium atom has a smaller nuclear charge

than neon and so a larger covalent radius

Fluorine’s greater nuclear charge pulls the outer energy

level in closer.

3+

-

radius = 134pm radius = 71pm

9+

-

--

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9+

-

--

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Atomic Size Summary

Decreasing Atomic Size

Across a period from left to right atomic size decreases

This is because of the atom having more electrons & protons and therefore a greater attraction which pulls the atom closer together hence the smaller size.

Atomic Size SummaryIn

crea

sing

Ato

mic

Siz

e

Down a group atomic size increases

This is because of the extra outer energy levels and the screening effect of the outer electrons.

Decreasing Atomic Size

Ionisation Energy

The ionisation energy is the energy required to remove one mole of electrons from one mole of atoms in the

gaseous state.

The first ionisation energy of magnesium:

Mg (g) Mg+ (g) + e- 744 kJmol-1

Values for ionisation energies can be found in the data book

Ionisation Energy

The third ionisation enthalpy shows a massive increase because it requires an electron to be removed from magnesium’s second energy level.

Mg2+ (g) Mg3+ (g) + e- 7750 kJmol-1

Mg+ (g) Mg2+ (g) + e- 1460 kJmol-1

The second ionisation energy of magnesium:

Looking across a periodFrom lithium to neon the first ionisation energy increases. Why?

Li (g) Li+ (g) + e- 526 kJmol-1

Ne (g) Ne+ (g) + e- 2090 kJmol-1

Li Be B C N O F Ne

An atom of Lithium

The lithium atom has 3 protons inside the nucleus

Li (g) Li+ (g) + e- 526 kJmol-1

3+

-

The outer electron is attracted by a relatively small nuclear charge

An atom of Neon

The neon atom has 10 protons inside the nucleus

10+

--

-

--

-

-

-

Ne (g) Ne+ (g) + e- 2090 kJmol-1

Each of neon’s eight outer electrons is attracted by a stronger nuclear charge

Looking down a group

The first ionisation energy decreases down a group in the periodic table. Why?

Li (g) Li+ (g) + e- 526 kJmol-1

Cs (g) Cs+ (g) + e- 382 kJmol-1

1. More Energy LevelsAs we saw with atomic size, the single electron in the outermost energy level is much further from the nucleus in caesium than in lithium.

Li

-

Caesium’s attraction for its outer electron is lowered by the screening effect caused by all its other electrons.

Cs

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2. Screening Effect

Ionisation Energy Summary

Increasing Ionisation Energy

Across a period from left to right ionisation energy increases

This is due to the increase in atomic charge having a greater pull on the electrons and therefore more energy is required to remove electrons.

Ionisation Energy SummaryD

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asin

g Io

nisa

tion

Ene

rgy

Down a group ionisation energy decreases

This is due to the outer electrons being further away from the nucleus and so the attraction is weaker and they are more easily removed.

Increasing Ionisation Energy

Atomic SizeThe atomic size is just like it sounds, the size of an atom

The covalent radius is defined as half the distance between the centres of covalently bonded atoms

The size of an atom will depend on two things:

The number of energy levels

The nuclear charge pulling the electrons in

Atomic SizeAcross a period from left to right the atomic size decreases. Down a group the atomic size increases from top to bottom.

Why?

Across a period an increasing nuclear charge (+) pulls the outer electrons (-) closer in towards the nucleus

Down a group there is an increase in the number of energy levels surrounding the nucleus

Ionisation EnthalpiesIonisation enthalpy is the quantity of energy required to remove 1 mole of electrons from 1 mole of atoms in the gaseous state.

ie. The energy required to remove an electron from an atom e.g.

Mg(g)

Mg+(g) + e-

What would affect the ionisation enthalpy?

How much attraction the outer electrons feel from the nucleus

Ionisation EnthalpiesAcross a period ionisation enthalpy increases from left to right

Down a group the first ionisation enthalpy decreases from top to bottom

What affects how much nuclear charge an electron feels?

Across a period increase in nuclear charge and decrease of atomic size makes it more difficult to remove an electron

Down a group there is a shielding (or screening) effect from the extra energy levels, and increased distance from the nucleus makes it easier to remove an electron

The first ionisation enthalpy is large for noble gases

There is a very large increase in ionisation enthalpy after the ion has achieved the noble gas configuration

Why?

It requires considerably more energy to remove an electron from a completely full shell (very stable) which is nearer the nucleus

Ionisation Enthalpies

Ionisation Enthalpies

The 1st ionisation energy is the energy needed to remove the first mole of electrons and the 2nd ionisation energy is the energy needed to remove the second mole of electrons, etc.

e.g. the ionisation energies for magnesium are:

1st Mg(g) à Mg+(g) + e- ∆H= 744 kJmol-12nd Mg+(g) à Mg2+(g) + e- ∆H= 1460 kJmol-13rd Mg2+(g) à Mg3+(g) + e- ∆H= 7750 kJmol-1

Why the big jump from 2nd to 3rd ionisation energy for Mg?

Ionisation EnthalpiesThere is a large jump in ionisation energy when the electron to be removed

comes from a new shell, closer to the nucleus.

ExamplesUse your Data Book to calculate the energy required for the following changes:

a) Ca(g) à Ca2+(g) + 2e-

b) Al(g) à Al3+(g) + 3e-

• The total energy to remove more than 1 mole of electrons is equal to the sum of each mole added together (as above).

Learning Outcomes :

Periodic Trends in Electronegativity05/09/2017

•I can explain the effect of electronegativity on electrons in a covalent bond.•I can describe, using a data book, the change in electronegativity down a group.• I can use atomic size and nuclear charge to explain the change in electronegativity across a period.•I can use atomic size and screening effect to explain the change in electronegativity down a group.

Electronegativity

Electronegativity is a measure of an atom’s attraction for the shared pair of electrons in a bond

e

eC H

Which atom would have a greater attraction for the electrons in this bond and why?

Linus PaulingLinus Pauling, an American chemist (and winner of two Nobel prizes!) came up with the concept of electronegativity in 1932 to help explain the nature of chemical bonds.

Today we still measure electronegativities of elements using the Pauling scale.

Since fluorine is the most electronegative element (has the greatest attraction for the bonding electrons) he assigned it a value and compared all other elements to fluorine.

Values for electronegativity can be found in the data book

ElectronegativitiesLooking across a row or down a group of the periodic table we can see a trend in values.

We can explain these trends by applying the same reasoning used for ionisation energies.

Looking across a periodIncreasing Electronegativity

Across a period electronegativity increasesThe charge in the nucleus increases across a period.Greater number of protons = Greater attraction for bonding electrons

What are the electronegativities of

these elements?

1.0

FCB N OLi Be

1.5 2.0 2.5 3.0 3.5 4.0

Looking down a group

4.0

3.0

2.8

2.6

F

Cl

Br

I

Dec

reas

ing

Ele

ctro

nega

tivity

Down a group electronegativity decreases

Atoms have a bigger radius (more electron shells)

The positive charge of the nucleus is further away from the bonding electrons and is shielded by the extra electron shells.

What are the electronegativities of

these halogens?

Attraction for bonding electronsDifferent elements have different attractions for bonded electrons

The relative ability of an element to attract electrons is called its Electronegativity

Those elements that require just one or two electrons to fill an energy level can attract electrons more easily

Electronegativity values increase across a period from LtoR

Electronegativity values decrease down a group

Therefore the most electronegative element is Fluorine

The smaller the atom the easier it is to capture an electron since they will feel a greater “pull” from the nucleus

Learning Outcomes :

Polar Covalent Bonds05/09/2017

•I can describe a covalent bond in terms of electrons and positive nuclei.•I can use electronegativity values to distinguish between pure covalent and polar covalent bonds.•I can use the data book to assign δ+ and δ- partial charges on atoms.•I can distinguish between bonds which show a dipole and those that are non-polar.

Experiment:

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Learning Outcomes :

Bonding Continuum05/09/2017

•I can explain the relationship between differences in electronegativity and type of bonding.•I can use data from properties such as conductivity, melting point and boiling point to deduce type of bonding and structure.•I can list exceptions to the statements: “Compounds formed between non-metals only are covalent. Compounds formed between a metal and a non-metal are ionic”.

Electronegativity & Bonding

The difference in the ability of elements to attract electrons tells us about the type of bonding we can expect between them

No difference = pure covalent bonding

Small difference (<1.5) = polar covalent bonding

Large difference (>1.5) = ionic bonding