They tend react mainly with non-metals to form ionic compounds witch are usually soluble white solids.
Untypical metallic properties:
- low melting points
- low density (first three float on water)
- very soft (easily squashed, extremely malleable)
- they have little material strength
Important trends down the group with increase in atomic number:
-
the melting point and boiling point generally decrease
-
Lithium melts at 180OC, Sodium at 98 OC, potassium at 64 OC, etc.
-
Lithium boils at 1347OC, Sodium at 883 OC, potassium at 774 OC, etc
- The elements gets more reactive
- … because the outer electron in more easily lost, because it is further from the nucleus
- They have to be stored in oil and handled with forceps (they burn skin).
- …because the outer electron is one extra full shell of electrons for each row we go down.
- Generally the density increases
- Although the atom gets bigger, there is greater proportional increase in the atomic mass
-
Lithium, Sodium and Potassium are all less dense than water. The others ‘float’ anyway, on H2
- Generally the hardness decreases
- Lithium is the hardest, but still easy to cut with a scalpel
- They are shiny when freshly cut, but soon go dull as they reacts with the air
Group 7 the Halogens (F, Cl, Br, I, At)
The elements of group 17, the Halogens, are a very similar set of mom-metals. They all exist as diatomic molecules (X2), and oxidize metals to form halides. The halogen oxides are acidic, and the hydrides are covalent. Fluorine is the most electronegative element at all. Generally, electronegativity and oxidizing ability decrease on descending the Group. Fluorine shows some anomalies because of the small size of its atom and ion.
The halogens are too reactive to occur free in nature. Fluorine is mined as fluorspar, calcium fluoride and cryolite. Chloride is also fund in minerals such as rock salt, and huge quantities of chloride ions occur in seawater, inland lakes subterranean brine wells. It is obtained by the electrolysis of molten sodium chloride or brine. Bromine is also found as the bromide ion in seawater, and in larger quantities in brine wells, from which it is extracted. Iodine is mined as sodium iodate (V), NaIO3 which is present in Chile saltpeter.
Physical features and important trends down the group with increasing atomic number:
- typical non-metals with relatively low melting and boiling points increase steadily down
-
the melting points increase steadily down the group (so the change in state room temperature form gas→liquid→solid ), this is because the intermolecular forces increase with increasing size of atom or molecule
- they are all colored non-metallic elements and the colour gets darker down the group
- they all have poor conductors of heat and electricity – typical of non-metals
- when solid they are brittle and crumbly e.g. iodine
- the size of the atom gets bigger as more inner electron shells are filled going down from one period to another, because there is less inclination to gain the extra full shell of electrons for each row we go down
Chemical features, similarities, physical property and reactivity trends:
- the atoms all have 7 outer electrons, this outer electron similarity, as with any Group in Periodic table, makes them have very similar properties e.g.
- They form singly charged negative ions e.g. Chloride (Cl) because they are one short of a noble gas electron structure. They gain one negative electron (reduction) to be stable and this gives a surplus electronic charge of -1. These ions are called the halide ions.
-
They form ionic compounds with metals e.g. sodium chloride Na+Cl -
- They form covalent compound with non-metals and with themselves. The bonding in the molecule involves single covalent bonds e.g. hydrogen chloride HCl or H – Cl
-
the elements all exist as X2 or X –X, diatomic molecules where X represents the halogen atom
- a more reactive halogen can displace a less reactive halogen from its salts.
- The reactivity decreases down the Group
Task 2
2.a
There more than 110 known elements. Atoms of these elements join together in different combinations to make countless millions of different compounds. Strong forces of attraction called chemical bonds hold atoms together in these compounds. Some compounds are very simple. For example, table salt contains just two elements, sodium and chlorine, bonded together. Its chemical name is sodium chloride. Other compounds are extremely complex, particularly substances found in living things, such as DNA and proteins.
Ionic bonding
Ionic bonding involves electron transfer between metals and non-metals. Both metals and non-metals try to achieve complete outer electron shell. Metals lose electrons from their outer shell and form positive ions (cations). This is an example of oxidation. Non-metals gain electrons into their outer shells and form negative ions. This is an example of reduction.
The strong electrostatic attraction between ions of opposite charge is an ionic bond. An ionic compound is composed of a giant regular structure of ions. This regular structure makes ionic compounds crystalline. The strong forces of attractions between ions make it difficult to separate the ions, and ionic compounds therefore have high melting and boiling points.
The ion formed by an element can be worked out from the position of the element in the periodic table. The elements in group 4 and group 8 (or 0) generally do not form ions.
The formula of an ionic compound: Ionic compounds consist of positive and negative ions. A sample of the compound is uncharged because the positive and negative charges balance exactly. In a sample of calcium chloride, CaCl2, the number of chloride ions, Cl-, is exactly twice the number of calcium ions, Ca2+2Cl- or CaCl2.
Covalent bonding
Covalent bonding involves electron sharing. It occurs between atoms of non-metals. It results in the formation of a molecule. The non-metal atoms try to achieve complete outer electron shells.
A single covalent bond is formed when two atoms each contribute one electron to a shared pair of electrons. For example, hydrogen gas exists as an H2 molecule. Each hydrogen atoms wants to fill its electron shell. They can do this by sharing electrons.
A single covalent bond can be represented by a single line. The formula of the molecule can be written as a displayed formula, H – H. The hydrogen and oxygen atoms in water are also held together by single covalent bonds.
Some molecules contain double covalent bonds. In carbon dioxide (CO2), the carbon atom has an electron arrangement of 2, 4 and needs and additional four electrons to complete its outer electron shell. It needs to share its four electrons with four electrons from oxygen atoms (electron arrangement 2, 6). Two oxygen atoms are needed, each sharing two electrons with the carbon atom.
Compounds containing covalent bonds have very different properties from compounds that contain ionic bonds. Covalent bonds are also strong bonds. They are intramolecular bonds- formed within each molecule. There are also intermolecular bonds- much weaker forces between the individual molecules.
The number of covalent bonds a non- metal atom can form is linked to its position in the periodic table. Metals (groups 1, 2, 3) do not form covalent bonds. The noble gases in group 8 are unreactive and usually do not form covalent bonds.
Metallic bonding
Metals are giant structures with high melting and boiling points.
Metal atoms give up one or more of their electrons for form positive ions, called cations. The electrons they give up form a ‘sea of electrons’ surrounding the positive metal ions and the negative electrons are attracted to the positive ions, holding the structure together.
The electrons are free to move through the whole structure, which is why metals conduct electricity. The electrons are delocalized, meaning they are not fixed in one position.
2.b
Ionic compound – sodium chloride
Ionic compounds form giant lattice structures. For example, when sodium chloride is formed by ionic bonding, the ions do not pair up. Each sodium ion is surrounded by six chloride ions, and each chloride ion is surrounded by six sodium ions.
The electrostatic attractions between the ions are very strong. The properties of sodium chloride can be explained using this model of its structure.
Covalent compound – Hydrogen
Covalent bonds are also strong bonds. They are intramolecular bonds- formed within each molecule. There are also intermolecular bonds- much weaker forces between the individual molecules. The properties of covalent compounds can be explained using a simple model involving these two types of bond or forces
Metallic bonding – potassium
Metallic bonds are different from ionic and covalent bonds. Some electrons from each metal atom are free to move from one atom to the next. The metal atoms become positive ions when they give up their electrons. They are held together in a lattice by a ‘sea’ of free electrons.
2c
The octet rule
The octet rule is a simple chemical rule of thumb that states that atoms tend to combine in such a way that they each have eight electrons in their valence shells, giving them the same electronic configuration as a noble gas. The rule is applicable to the main-group elements, especially carbon, nitrogen, oxygen, and the halogens, but also to metals such as sodium or magnesium. In simple terms, molecules or ions tend to be most stable when the outermost electron shells of their constituent atoms contain eight electrons.
Ionic bonding
The ionic bond is the strong electrical attraction between the positive and negative ions next to each other in the lattice. These strong bonding forces makes the structure hard (if brittle) and have high melting and boiling points.
Ionic bonding involves electron transfer between metals and non-metals. Positive ions are formed when electrons are removed from atoms. This happens most easily with metallic elements. Atoms of non-metallic elements tend to gain electrons to form negative ions. When metals combine with non-metals, electrons are transferred from the metal atoms to non-metals atoms. Usually a metal atom will lose all of its outer-shell electrons and a non-metal atom will accept electrons to fill its outer shell. The net result of electros transfer from a metal atom to a non-metal atom to produce filled outer shell similar to the noble-gas
Task 3
3a
Exothermic Reaction
Theory 1:
If the products contain less energy than the reactants, heat is released or given out to the surroundings and the change is called exothermic.
The reaction between hydrochloric acid (HCl) and sodium hydroxide (NaOH) is exothermic and can be represented by the equation:
HCl + NaOH → NaCl + H2O
Chemicals & apparatus:
20g 2M NaOH solution;
20g 2M HCl solution;
Thermometer
Polystyrene cup
Beaker
Measuring cylinder
Method:
I measured 20.0 cm3 of 2M HCl and poured into the polystyrene cup. I noted the temperature which was 21oC. Then I measured 20.0 cm3 of 2M NaOH. The temperature of that alkali was 22oC. Next, I poured it into the beaker containing the acid. Then, I recorded the maximum temperature, which was 32oC.
Results:
Theory 2:
If the specific heat capacity of the solution is c and its mass is m then the heat required to release its temperature by ΔT is:
Q (J) = m x c x (ΔT)
If we know the mass of solution, its c-value (specific heat capacity) and the number of moles of reacting H+ and OH- giving a temperature rise of ΔT, the enthalpy of the reaction can be found.
Calculations 1:
20.0 cm3 of HCl – initial temperature = 21oC final temp of the compound=32oC
20.0 cm3 of NaOH – initial temperature = 22oC
Average initial temperature: 22+21
2 = 21.5oC
Rise in temp, ΔT = 32 – 21.5 = 10.5
Mass= 20.0 g + 20.0 g= 40 g, because 1 cm3 = 1g
Heat (Q) = m x c x (ΔT)
= 40 x 4.2 x 10.5
= 1764J
= 1.764 kJ
It means that 1.746 kJ is given out when 20 cm3 of acid (HCl) mixed with 20 cm3 of alkali (NaOH).
Calculations 2:
The equation for the neutralization is: HCl + NaOH → NaCl + H2O.
Since the solutions are both 2M, 20cm3 of the each contain 20 x 2 /1000 mole of H+ and OH-
H+ + OH- →H2O ΔH= -57.3kJmol-1
Acid = 2M = 2 mole per 1000cm3
Alkali = 2M = 2 mole per 1000cm3
Endothermic Reaction
Theory 1:
………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………
NH4Cl (s) + H2O (l)→ NH4OH + HCl
Chemicals & apparatus:
3g of NH4Cl (ammonium chloride)
50g of H2O (water)
Thermometer
Polystyrene cup
Beaker
Measuring cylinder
Laboratory balance
Method:
I measured 50.0 cm3 (= 50g) of H2O and poured into the polystyrene cup. I noted the temperature which was 23oC. Then I measured 3g of NH4Cl. Next, I added it into the polystyrene cup which was containing the water. Then, I recorded the final temperature which was 18oC.
Results:
Theory 2:
If the specific heat capacity of the solution is c and its mass is m then the heat required to release its temperature by ΔT is:
Q (J) = m x c x (ΔT)
Calculations 1:
50.0 cm3 of H2O – initial temperature = 23oC final temp of the compound=18oC
3g of NH4Cl – solid
Initial temperature = 23oC
Final temperature= 18oC
Rise in temp, ΔT = 23 – 18 = 5oC
Mass= 50.0 g + 3.0 g= 53 g, because 1 cm3 = 1g
Heat (Q) = m x c x (ΔT)
= 53 x 4.2 x 5
= 1113J
= 1.113 kJ
It means that 1.113 kJ is involved when 50 cm3 of H2O mixed with 3g of NH4Cl.
Calculations 2:
NH4Cl – 1 mole
Relative atomic mass of NH4Cl: N=14; H=1; Cl=35.5
N+4xH+Cl
14+ (4x1) +35.5 = 53.5
53.5g = 1 mole
3g =3/53.5 mole ≈0.056 mole
3b
The heat change is called the enthalpy change is denoted by delta H, ΔH.
- ΔH is negative (-ve) for exothermic reactions i.e. heat energy is given out and lost from the system to the surroundings which warm up.
- ΔH is positive (+ve) for endothermic reactions i.e. heat energy is gained by the system and taken in from the surroundings which cool down OR, as is more likely, the system is heated to provide the energy needed to effect the change.