periodic table of elements task 1a (1)

8/4/2019 Periodic Table of Elements Task 1a (1)

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1 Task 1a – Unit 1 Fundamentals in Science (BTEC National Diploma in Applied Sciences)

Otgon Orgil

Periodic Table of Elements

Many of the elements (in compounds and their solutions) involved in carrying out the

preparation of standard solutions and titrations are in different groups and periods of

the periodic table.Using the elements sodium, potassium, calcium, magnesium, carbon, oxygen,

nitrogen, fluorine, chlorine, iron, hydrogen and sulphur, provide some key features of

the periodic table such as:

Their atomic number and relative atomic masses

Their electronic structures

How they are organised in groups, periods and blocks.

Any trends in their chemical properties (chemical reactivity down group I with

water & displacement reactions down group VII). Use balance chemical

equations to explain the trend down group I and VIII

Any trends in physical properties (Ionisation Energy & atomic radii) down a

group and across a period

Atomic number: 11

Relative Atomic Mass: 22.98976928 (2)

Group in periodic table: 1

Group name: Alkali metal

Period in periodic table: 3

Block in periodic table: s-block

Ionization energy (eV): 5.1391

Calculated Atomic Radii: 190

Electronic configurations: 2,8,1 or [Ne] 3s1

Atomic number: 19








Electronic configurations: 2,8,8,1 or [Ar] 4s1

Atomic number: 20








Electronic configurations: 2,8,8,2 or [Ar] 4s2

11

Na Sodium

22.99

19

K Potassium

39.10

20

Ca Calcium

40.08




Otgon Orgil

Atomic number: 12



Group name: Alkaline earth metal





Electronic configurations: 2,8,2 or [Ne] 3s2

Atomic number: 6Relative Atomic Mass: 12.0107 (8)


Group name: NA


Block in periodic table: p-block



Electronic configurations: 2,4 or [He] 2s2 2p2

Atomic number: 8



Group name: Chalcogen






Atomic number: 7



Group name: Pnictogen






12

Mg Magnesium

24.31

6

C Carbon

12.01

8

O Oxygen

16.00

7

N Nitrogen

14.01




Otgon Orgil

Atomic number: 9



Group name: Halogen






Atomic number: 17

Relative Atomic Mass: 35.453 (2)Group in periodic table: 17

Group name: Halogen





Electronic configurations: 2,8,7 or [Ne] 3s2 3p5

Atomic number: 26



Group name: NA


Block in periodic table: -block



Electronic configurations: 2,8,14,2 or [He] 3d6 4s2

Atomic number: 1Relative Atomic Mass: 1.00794 (7)


Group name: NA





Electronic configurations: 2,7 or 1s1

9

F Fluorine

19.00

17

Cl Chlorine

35.45

26

Fe Iron

55.85

1

H Hydrogen

1.01




Otgon Orgil

Atomic number: 16



Group name: ChalcogenPeriod in periodic table: 3




Electronic configurations: 2,8,6 or 3s2 3p4

Trends in chemical properties

I have observed seven (7) of the elements of the above are in the reactivity series of

metals. K, Na and Ca react with water in displacement. Mg, Fe and C react with

acids.

Going from bottom to top, the metals:

increase in reactivity;

lose electrons more readily to form positive ions; corrode or tarnish more readily;

require more energy (and different methods) to be separated from their ores;

become stronger reducing agents.

I have also observed group VII (Halogens) have water and displacement reactions.

The halogens become less reactive down the group. So bromine is less reactive thanchlorine, and iodine is less reactive than bromine. A more reactive halogen will

displace a less reactive halogen from an aqueous solution of its ions. Moreover, the

boiling points of the elements increase, elements become darker and less reactive

as oxidising agents when going down the group. They will also react with Group I to

form salts.

16

S Sulfur

32.07




Otgon Orgil

Definitions

Atomic mass (ma) is the total mass of protons, neutrons and electrons in a singleatom. There are three types of atomic masses: relative atomic mass, standard

atomic weight and relative isotopic mass. In this assignment, I will be predominately

using relative atomic mass primarily found in the periodic table. The relative atomic

mass is closest to the average atomic mass found in a particular sample, weighted

by isotopic abundance.

Atomic number (also known as the proton number) is the amount of protons found in

the nucleus of an atom, identical to the charge number of the nucleus. It is

conventionally represented as Z. The atomic number uniquely identifies a chemical

element.

A period is elements that are arrange in a series of rows so that those with similar

properties appear in a vertical column. Elements of the same period have the same

number of electron shells; with each group across a period, the elements have one

more proton and electron and become less metallic. This arrangement reflects the

periodic recurrence of similar properties as the atomic number increases. For

example, the alkaline metals lie in one group (group 1) and share similar properties,

such as high reactivity and the tendency to lose one electron to arrive at a noble-

gas electronic configuration.

A block in a periodic table is a set of adjacent groups. Each block is named after its

characteristic orbital.

There are 4 primary blocks and 1 hypothetical:

• s-block (alkali meals & alkaline earth metals - groups 1-2, periods 1-7)

• p-block (nonmetals & semimetals as well as some metals groups 13-18, periods 2-7)

• d-block (transition metals - groups 3-12, periods 4-7)

• f-block (inner transition metals – Lanthanoids plus Actinoids)

The following is the order for filling the "subshell" orbitals which also gives the linear

order of the "blocks" (as atomic number increases) in the periodic table: 1s, 2s, 2p, 3s,

3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.




Otgon Orgil

A group in a periodic table is known as a family, a vertical column of chemical

elements. There are 18 groups in the periodic table. The modern explanation of thepattern of the table is that the elements in a group have similar configurations of the

outmost electron shells of the atoms (stabilization and destabilization, the same as

the core charge).

New IUPAC numbering Name

Group 1 the alkali metals or lithium family

Group 2 the alkaline earth metals or beryllium family

Group 3 the scandium family

Group 4 the titanium family

Group 5 the vanadium family

Group 6 the chromium family

Group 7 the manganese family

Group 8 the iron family

Group 9 the cobalt family

Group 10 the nickel family

Group 11 the copper family

Group 12 the zinc familyGroup 13 the boron group or boron family

Group 14 the carbon group or carbon family

Group 15 the pnictogens or nitrogen family

Group 16 the chalcogens or oxygen family

Group 17 the halogens or fluorine family

Group 18 the noble gases or helium family or neon family

Atomic radius is a term used to describe the

size of the atom, but there is no standard

definition for this value. Atomic radius may

refer to the ionic radius, covalent radius,metallic radius, or van der Waals radius. In all

cases, the size of the atom is dependent on

how far out the electrons exte nd. The

atomic radius for an element tends to

increase as one goes down an elementgroup. The electrons become more tightly

packed as you move across the periodic

table, so while there are more electrons for

elements of increasing atomic number, the atomic radius actually may decrease.

Group IAElements

ElectronicConfigurations

Numberof Shells

Group IIAElements

ElectronicConfigurations

Numberof Shells

Hydrogen 1 1 Beryllium 2,2 2

Lithium 2,1 2 Magnesium 2,8,2 3

Sodium 2,8,1 3 Calcium 2,8,8,2 4

Potassium 2,8,8,1 4 Strontium 2,8,18,8,2 5

Rubidium 2,8,18,8,1 5 Barium 2,8,18,18,8,2 6

Cesium 2,8,18,18,8,1 6 Radium 2,8,18,32,18,8,2 7

Francium 2,8,18,32,18,8,1 7




Otgon Orgil

Ionization energy is the amount of energy required to remove the highest-energy

electron from an isolated neutral atom in the gaseous state. Ionization Energy haspositive values because energy is always required to remove an electron, it is

endothermic. Electrons are attracted to the nucleus therefore energy is needed to

remove them.

Ionization Energy increases when: there is stronger nuclear charge - there is a full

subshell (therefore the atom is stable) - there is a half full subshell therefore there is

evenly distributed charge in the atom.

Ionization Energy is low when: there is more shielding - the electron to be removed is

spin paired (therefore there is repulsion) - the electron is far from the nucleus - when

an electron is in the next subshell, Ionization Energy will lower because the outer

electron is at a greater distance from the nucleus and there is more shielding from

inner subshells - when there is a half full subshell in an atom, the first Ionization Energy

is high because the atom is stable.

Tips

For any element:

Number of Protons = Atomic Number

Number of Electrons = Number of Protons = Atomic Number

Number of Neutrons = Mass Number - Atomic Number

Remember PEN

periodic table of elements task 1a (1)

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