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GCSE Chemistry > C3 - Structure and Bonding > Flashcards

C3 - Structure and Bonding Flashcards

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1
Q

What is Ionic Bonding?

A

The transfer of electrons between a metal and a non-metal.

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2
Q

What is an Ionic Bond?

A

The strong bond of oppositely charged ions that are held together by strong electrostatic forces of attraction.

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3
Q

What holds ions together?

A

Strong electrostatic forces of attraction between oppositely charged ions.

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4
Q

What ions do metals form in an Ionic Compound?

A

Positive Ions - lose electrons

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5
Q

What ions do non-metals form in an Ionic Compound?

A

Negative Ions - gain electrons

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6
Q

How are the ionic bonds arranged?

A

In giant structures/lattices

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7
Q

What are the charges on the ions for each group?

A

Group:

1) 1+
2) 2+
3) 3+
4) 4+ or 4-
5) 3-
6) 2-
7) 1-
0) never as they’re electronically stable

Rule:

  • Group 1 atoms have 1 electron in their outermost shell.
  • Group 2 atoms have 2 electrons in their outermost shell.
    etc. ……
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8
Q

What is the ‘Dot and Cross’ Method and how does it help with understanding Ionic Bonding?

A

The ‘Dot and Cross’ Method helps understand what occurs when atoms from 2 elements form an Ionic Bond. It makes it easier to visualise the transfer of electrons.

(See GoodNotes 3.3)

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9
Q

What happens when Mg (2.8.2) forms an Ionic Bond with O (2.6)?

A

Mg (2.8.2) reacts with O (2.6) to form MgO.

Mg needs to lose 2 electrons. O needs to gain 2 electrons also. To become an electronically stable compound, Mg transfers 2 electrons to the outershell of O. Magnesium Oxide (MgO) is formed (2.8).

Within the ionic compound:

  • Mg (2.8.2) beomces [Mg]2+ (2.8)
  • O (2.6) becomes [O]2- (2.8).
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10
Q

What happens when Al forms an Ionic bond with F?

A

Al (2.8.3) reacts with F (2.7) to form AlF3.

Aluminium needs to lose 3 electrons. But, Fluorine only needs to gain 1. So, you need 3 times as many fluorine atoms to each receive 1 electron to become stable.

Within the ionic compound:

  • Al (2.8.3) becomes [Al]3+ (2.8)
  • F (2.7) beomes [F]1- (2.8) - there are three [F]1- ions in the compound.
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11
Q

What are the chemical properties of a Giant Ionic Structure?

A
  • High Melting Point - strong ionic bonds operating in all directions.
  • Solid at room temperature
  • Great conductor of electricity and heat - when in a molten or dissolved state!!!
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12
Q

When do ionic compounds conduct electricity?

A

Once in a molten or dissolved state:

  • the ions are free to move around and are attracted to oppositely charged ions.
  • they can carry the charge through the liquid.
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13
Q

Why can ionic compounds only conduct electricity when in a molten or dissolved state?

A

The ions must be mobile (free to move around).

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14
Q

Why do ionic compounds NOT conduct electricity when in a solid state?

A

The ions are rigidly fixed in place - they are immobile and cannot carry a charge.

Most ionic compounds are solid at room temperature and therefore have to be melted or dissolved in water.

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15
Q

What electrode are metal and non-metal ions attracted to?

A

Metal ions - positively charged —> NEGATIVE ELECTRODE

Non-metal ions - negatively charged —> POSITIVE ELECTRODE

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16
Q

Why can seawater conduct electricity better than fresh water?

A

Because sea water has lots of ions (sodium ions and chloride ions). Therefore, an electrical charge is carried through seawater more easily than in freshwater.

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17
Q

What is Covalent Bonding?

A

The sharing of PAIRS of electrons in NON-METALS to form molecules.

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18
Q

What is a Covalent Bond?

A

The strong bond of atoms held together by shared pairs of electrons.

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19
Q

What are Intramolecular bonds?

A

The bonds within a molecule (between atoms).

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20
Q

What are Intermolecular bonds?

A

Bonds between the molecules of a substance.

This impacts melting point and boiling point.

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21
Q

Which 2 ways can you represent covalent bonding visually?

A

1) Venn Diagram with Dot and Cross
2) Number or covalent bonds e.g:

H—O—H

-Make sure it’s === for double, and — for single

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22
Q

How do 2 Hydrogen atoms form a covalent bond?

A

Both hydrogen atoms have 1 electron in their outermost shells as they are of the same element. Therefore they need to lose 1 electron to become stable. This means 1 pair of shared electrons is needed which forms a single covalent bond.

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23
Q

How do 2 Oxygen atoms form a covalent bond?

A

They both need to lose 2 electrons as oxygen is in group 6. Therefore 2 shared pairs of electrons are needed which means that a DOUBLE covalent bond is formed.

It’s a double as there are 2 shared pairs in 1 intersect in imaginary venn diagram.

24
Q

How does a Hydrogen atom form a covalent bond with a Chlorine atom?

A

Hydrogen needs to lose 1 electron and chlorine needs to gain 1 so 1 shared pair is needed to form a single covalent bond.

If you had to show this in a ‘Dot and Cross’, work out how many shared pairs there are and then fill in the rest of the electrons that are missing.

25
Q

How do 2 hydrogen atoms form a covalent bond with an oxygen atom? (Water)

A

There are 2 hydrogen atoms which means that they each want to lose 1 electrons. That’s 2 electrons that need to be lost in total. Oxygen wants to gain 2 electrons. Therefore 2 shared pairs of electrons are needed and water has 2 SINGLE covalent bonds. NOT A DOUBLE.

They are single because there are 2 intersects in the imaginary venn diagram.

26
Q

What are the 2 types of Covalently-bonded substances?

A

1) Simple covalent structures - made up of MOLECULES. (E.g H2O, CO2 etc…)
2) Giant covalent structures - made up of complex three-dimensional structures of ATOMS. (E.g Diamond, Graphite, Silicon Dioxide)

27
Q

What are the intramolecular and intermolecular forces like for simple covalent structures? What does this mean?

A

-Consist of molecules in which the atoms are joined by strong covalent bonds - STRONG INTRAMOLECULAR BONDS.

WEAK INTERMOLECULAR FORCES between the molecules - low melting points and low boiling points.

28
Q

In a simple covalent structure, the Intermolecular forces are ______.

A

Weak - covalent substances with simple structures are easier to break down (lower melting and boiling points).

29
Q

In a simple molecule, the Intramolecular forces are ______ because…..

A

Strong because of the strong covalent bonds that make up each atom.

30
Q

Can simple covalent structure conduct electricity?

A

No. The neutral molecules cannot conduct electricity at all.

31
Q

What are giant covalent structures? How are they arranged?

A

Complex three-dimensional network of atoms held together by strong covalent bonds.

Arranged in giant, regularly-structured lattices which are extremely strong due to the many bonds involved.

32
Q

What are 3 substances that are made up of giant covalent structures?

A

1) Diamond
2) Graphite
3) Silicon Dioxide

33
Q

What special properties do substances made up of giant covalent structures have?

A
  • Very high melting points and boiling points.
  • Insoluble in water
  • Extremely Hard
  • DO NOT conduct electricity (except graphite!!!)
34
Q

What is the hardest natural substance known to man? What is its boiling point?

A

Diamond

BP = 4827°C

35
Q

Why is diamond so hard?

A

Each carbon atom forms 4 strong covalent bonds, arranged in a perfectly symettrical lattice.

In order to break a diamond, immense pressure is needed.

36
Q

How are the atoms in graphite (form of carbon) arranged?

A
  • The carbon atoms are arranged in hexagons which form different layers.
  • Between the layers, there are no covalent bonds, so the layers can EASILY SLIDE OVER EACH OTHER. (*Like cards sliding off the deck of a pack of playing cards.)
  • Each carbon atom forms three strong covalent bonds.
  • But carbon has 4 electrons in its outer shell - leaving one electron per atom spare.
37
Q

What are delocalised electrons?

A

Delocalised electrons are electrons that can move freely around in a substance.

38
Q

What role do delocalised electrons play in graphite?

A

Delocalised electrons enable graphite to be able to conduct electricity as each mobile outer electron can move freely along the layers of carbon atoms.

39
Q

What are the properties of graphite?

A

1) Excellent electrical and *thermal conductor -
* each delocalised electron gains energy and moves faster - rapidly transferring thermal energy along the layers of graphite.
2) Smooth, slippery and soft.

40
Q

Why can diamond not conduct electricity?

A

Because diamond has no delocalised electrons.

Graphite is the ONLY Giant Covalent Structure that can conduct electricity.

41
Q

What is graphite used as?

A

As pencil lead and a lubricant.

42
Q

What are fullerenes? What shape are they?

A

Fullerenes are hollow, spherical, cage-like carbon molecules, usually made up of *hexagonal rings of carbon atoms. (E.g C50, C60, C70)

Shape = SPHERICAL

*Sometimes made up of pentagonal and heptagonal rings of carbon.

43
Q

When and by whom was it discovered that carbon atoms could behave like this?

A

1985 - Scientists scientists tried to re-create conditions that accounted for long chains of carbon atoms that existed in outer space.

By doing this, they created the C60 molecule by chance - it had 60 carbon atoms - named ‘Buckminsterfullerene.

Professor Sir Harry Kroto suggested that the carbon atoms were arranged in a cage-like structure, similar to the panels of a football.

44
Q

What are the properties of fullurenes? What role could fullerenes play in the world of medicine in the future?

A

Properties:

  • High Tensile Strength
  • High electrical and thermal conductivity (delocalised electrons)

Future (medical industry):
-Fullerenes could be used for drug-delivery into the body. For example, the cage-like structures could be used as ‘bucky mules’ to deliver radioactive atoms to fight cancer.

45
Q

What is graphene? How was it formed?

A

Graphene is an exciting new material that was made when scientists at Manchester University used a piece of tape to rip off a layer of graphite. They then looked at it under a powerful electronic microscope.

Summary:
-Graphene is a single layer of graphite and therefore is one atom thick.

46
Q

What are the properties of graphene?

A
  • Excellent electrical and thermal conductor (even better than graphite).
  • Low density
  • Very reactive - most reactive form of carbon.
47
Q

What could graphene’s first large-scale application be?

A

Graphene’s first large scale application is likely to be as flexible electronic displays - watching a film on your wrist or sleeve.

48
Q

How are the atoms in a metal arranged?

A

The atoms in metals are closely packed and are arranged in regular layers, one on top of another.

49
Q

How does metallic bonding work? What properties does this give metals?

A

Metals are arranged in giant lattices of positively charged ions.

  • When atoms are arranged like this in a metal, they all give up their outer electrons to share them in the lattice.
  • This makes all of the metal atoms in the lattice positive ions.
  • The outer electrons shared throughout the metals are called delocalised electrons. These electrons can move freely throughout the lattice, forming a ‘sea of delocalised electrons’.
  • The positively charged metal ions are held together by strong forces of attraction between them and the negative electrons.

Properties:

  • Very strong - regular structure
  • High melting and boiling points - regular structure
  • Good conductor of electricity and heat - delocalised electrons
50
Q

Why can pure metals be bent and shaped?

A

Pure metals are very malleable and can be bent and shaped.

This is because the regular layers of atoms in a pure metal are able to slide over each other easily.

51
Q

Why can the layers of pure metals slide over each other easily?

A

Because of the regular arrangement of layers within a giant metallic structure.

52
Q

What is an alloy?

A

An alloy is a mixture of 2 or more elements, at least one of which is a metal.

53
Q

Why are alloys harder than pure metals?

A

Because the atoms of different elements distort the regular arrangement of a metal - so the layers can no longer slide over eachother.

54
Q

Why can thermal and electrical energy be transferred through metals?

A

Because of the ‘sea’ of delocalised electrons that can carry thermal and electrical energy through the metallic structure.

55
Q

Why do metals have high melting points?

A

Because of the regular and structured layers held together by strong electrostatic forces of attraction.

GCSE Chemistry (17 decks)

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