PERIODIC TABLE Flashcards
Can you explain the basis for the arrangement of elements in the periodic table?
Elements in the periodic table shows periodicity. Periodicity is the recurrence of similar chemical and physical properties at regular intervals.
Elements are organized on the basis of
- increasing atomic number
- electronic configuration
- chemical properties
Origins of the Periodic Table?
Johann Dobereiner (proposed the law of triads) He noticed that certain groups of three elements, which he called triads, showed similar chemical and physical properties. The elements in any triad were arranged in increasing relative atomic mass and the relative atomic mass of the middle element was close to the average of the first and third elements. Example lithium, sodium and potassium have relative atomic masses of 7, 23 and 39.
John Newlands (proposed the law of octaves) Arranged the elements that had been discovered at the time in order of increasing relative atomic mass and found that each element exhibited similar chemical and physical properties to the element eight places ahead of it in the list. Example, sodium was eight places ahead of lithium and the two exhibited similar properties. He then placed the similar elements into vertical columns called groups.
Dimitri Mendeleev
Arranged elements by increasing relative atomic mass. Placed elements with similar chemical and physical properties together in vertical columns (groups). Occasionally ignored the order suggested by relative atomic mass and exchanged adjacent elements so they were better classified into chemical families. Mendeleev is credited with the first version of the periodic table.
Henry Moseley
Placed the elements in increasing atomic number which resulted in all elements with similar properties falling in the same group.
How can electronic configuration be calculated from the periodic table?
If the group number and period number of an atom are known, its electronic configuration can be determined. And if the electronic configuration of an atom is known, the group number and period number can be determined.
The group number and the valence number are the same. The period number and the number of occupied electron shells are the same.
Potassium is in Group 1 and Period 4
1 valence electron, 4 occupied electron shells
E.C of a potassium atom is 2, 8, 8, 1
E.C of a silicon atom is 2, 8, 4
4 valence electrons, 3 occupied electron shells
Silicon is in Group IV and Period 3
Can you explain trends in group II?
Alkaline earth metals
Cations
Elements in Group II all have similar chemical properties because their atoms all have two valence electrons. They react by losing these valence electrons to form positively charged ions called cations. When they lose these electrons they are said to ionise. The easier an element ionises, the more reactive it is. The ease of ionisation increases moving down Group II, therefore the reactivity of the elements increases moving down the group.
— down Group 2
- ease of ionization increases
- reactivity of the elements increases
- atomic radius increases
Beryllium number of occupied e- shells = 2
Magnesium = 3
Ca = 4
- Atomic Radius = Increases moving down due to the increase in number of occupied electron shells.
- Ease of Ionization = Increases moving down. As the atomic radii increase, the attractive pull of the positive nucleus on the valence electrons decreases, and the more easily the atoms lose their valence electrons to form cations.
Reactivity with oxygen, water and dilute HCl = Increases moving down due to the increase in the ease of ionization.
Examples with calcium
2Ca(s) + O2(g) = 2CaO(s)
Ca(s) + 2H2O(l) = Ca(OH)2(aq) + H2(g)
Ca(s) + 2HCl(aq) = CaCl2(aq) + H2(g)
Can you explain trends in group VII?
Halogens
Anions
Elements in Group VII exist as diatomic molecules, these being F2 Cl2 Br2 and I2. They all have similar chemical properties because their atoms all have seven valence electrons. They react by gaining one valence electron to form negatively charged ions called anions. When they gain this electron they are said to ionise. The ease of ionisation increases moving up Group VII, therefore the reactivity of the elements increases moving up the group.
— up
- ease of ionization increases
- reactivity of the elements increases
- strength of oxidizing power increases
F
Appearance and state at room temp = pale yellow gas
Number of occupied e- shells = 2
Cl
Appearance and state at room temp = yellow-green gas
Number of occupied e- shells = 3
- Atomic radius = Increases moving down due to the increase in number of occupied electron shells.
- Ease of Ionization = Increases moving up. As the atomic radii decrease, the attractive pull of the positive nucleus on the electron to be gained increases and the more easily the atom gains this electron to form anions.
- Reactivity = Increases moving up due to the increase in ease of ionization.
- Strength of oxidizing power = Increases moving up due to the increase in ease of ionization. The more easily the element ionizes, the more easily it takes electrons from another reactant. Chlorine will take electrons from bromide (Br-) and iodide (1-) ions and bromine will take electrons from iodide (1-) ions.
- Displacement = An element is displaced from its compounds by an element above it in the group. (A more reactive element will displace a less reactive element. Chlorine will displace bromine and iodine, and bromine will displace iodine from their compounds).
In a displacement reaction an element in its free state takes the place of another element in a compound. A more reactive element will displace a less reactive element. Chlorine will displace bromine and iodine, and bromine will displace iodine from their compounds.
CI2(g) + 2KBr(aq) = 2KCI(aq) + Br2(aq)
Displacement reactions can be explained by looking at the relative strength of oxidising power of the elements. This is determined by how eaSily one substance takes electrons from another substance. The strength of oxidising power of Group VII elements increases moving up the group because the ability to ionise and take electrons from another reactant increases moving upwards. Chlorine will take electrons from bromide (Br-) and iodide (1-) ions and bromine will take electrons from iodide (1-) ions.
CI2(g) + 2Br-(aq) = 2C1-(aq) + Br2(aq)
Can you explain trends in period 3?
Moving along Period 3 from left to right, the metallic nature of the elements decreases and the non-metallic nature increases. Each element has three occupied electron shells.
Na = 2,8,1 (Metal) (Good Conductors)(Loses 1e-) Mg = 2,8,2 (Metal) (Good Conductors) (Loses 2e-) Al = 2,8,3 (Metal) (Good Conductors) (Loses 3e-)
Ease of ionization = Increases moving from right to left so decreases moving along the period. As the atomic radii increase and the number of positive protons decreases, the more easily the atoms lose electrons to form positive cations. Reactivity = increases moving from right to left due to the increase in ease of ionisation.
Si in Group IV = 2,8,4 (Semi-Metal/Metaloid)
(Semi-Conductor) (Shares 4e-)
Does not usually ionize and reacts by sharing e- with other non-metal atoms.
P = 2,8,5 (Non-Metal) (Insulator) (Gains 3e-) S = 2,8,6 (Non-Metal) (Insulator) (Gains 2e-) Cl = 2,8,7 (Non-Metal) (Insulator) (Gains 1e-)
Ease of ionization = Increases moving from left to right/along the period. As the atomic radii decrease and the number of positive protons increases, the more easily the atoms gain electrons to form negative anions. Reactivity = increases moving from left to right due to the increase in ease of ionisation.
Ar = 2,8,8 (Non-Metal) (Insulator) (None)
Does not ionise
Chemically Unreactive
Atomic radius = decreases moving from left to right due to the increase in number of positive protons causing the attractive pull of the positive nucleus on the valence electrons to get stronger.
- The ease of ionisation and reactivity of the metals sodium, magnesium and aluminium decreases moving along the period.
- The ease of ionisation and reactivity of the non-metals phosphorus, sulfur and chlorine increases moving along the period.
- Silicon does not usually ionise, it usually reacts by sharing electrons with other non-metal atoms.
- Argon does not ionise and is chemically unreactive.
— across
- from left to right the metallic nature of the elements decreases and the non-metallic nature increases
- ease of ionization and reactivity of sodium, magnesium and aluminium decreases
- ease of ionization and reactivity of non-metals phosphorus, sulfur and chlorine decreases
- atomic radius decreases moving from left to right
Group 0?
When non-metal atoms with 4, 5, 6 or 7 valence electrons bond with each other, they share valence electrons and form molecules.
Notes
Metals are found on the left side of the table and non-metals are found on the right side of the table.
Groups
- down
- each element has one more electron shell than the element directly above it
- the metallic nature of elements increases and the non-metallic nature decreases
For elements in Groups I to VII, the number of valence electrons is the same as the group number. All elements in Group 0 have a full outer electron shell. Moving down any group, each element has one more electron shell than the element directly above it. All elements in the same group have similar chemical properties. Moving down a group, the metallic nature of elements increases and the non-metallic nature decreases.
Between Groups II and III there are ten groups of elements called the transition elements or transition metals. Transition meta ls usually have two valence electrons.
Periods
— across
* from left to right each element has one more valence electron than the element directly before it
* from left to right the metallic nature of the element decreases and the non-metallic nature increases
All the elements in the same period have the same number of occupied electron shells, therefore
they have their valence electrons in the same shell. The number of occupied electron shells is the same as the period number.
Moving along any period from left to right, each element has one more valence electron than the element directly before it. Moving along any period from left to right, the metallic nature of the elements decreases and the non-metallic nature increases.
What are groups in the periodic table?
Groups are vertical columns of elements.
What are periods in the periodic table?
Periods are horizontal row of elements.
