Unit 1 – (B) Periodicity subtopic assessment Feed back – General thoughts • On the whole progress is evident when compared to Unit 1 - (A)

Controlling rate.• Most on target or just below it, discuss targets today or tomorrow

and review progress.

• Terminology, description, definition and answer structure has areas for improvement

• Open-ended questions need a lot of practice.

Multiple choice • Spend 5-10 minutes reviewing answers, identify questions that posed you trouble and

discuss them in your group. I will be circulating the classroom to deal with any concerns raised or points you wish to discuss.

1) B 11) A 21) C2) D 12) C 22) N/A3) C 13) N/A 23) D4) A 14) B 24) D5) B 15) A6) B 16) D7)D 17) D8) C 18) C9) C 19) N/A10) D 20) C

EXTENDED RESPONSEQ1) …Clearly explain why the first ionisation energy decreases down this group. (2)

• Distance from nucleus and valence electron increases as you go down group 1, electrostatic interaction decreases as a result. (1)

• Electron Shielding by core electrons increases down group 1, valence electron requires less energy to remove from atom. (1)

Electron shielding: The shielding effect describes the attraction between an electron and the nucleus in any atom with more than one electron shell. Shielding effect can be defined as a reduction in the effective nuclear charge on the electron cloud, due to a difference in the attraction forces of the electrons on the nucleus.

EXTENDED RESPONSE• Q2) a) what is the trend in electronegativity values across this period from

Li to F. (1)

It increases. (1)

Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. The Pauling scale is the most commonly used. Fluorine (the most electronegative element) is assigned a value of 4.0, and values range down to caesium and francium which are the least electronegative at 0.7.

EXTENDED RESPONSE• Q2) b) ….Why is the fourth ionisation energy of aluminium so much higher

than the third ionisation energy? (1)

• The fourth ionisation would require additional energy because the electron is being removed from an inner shell, which is closer to the nucleus. (1)

Ionisation energy: The energy required to remove one mole of electrons from one mole of atoms in the gaseous state.

EXTENDED RESPONSE• 2) c) ….Why do the boiling points increase down group 7? (1)

• London dispersion forces are stronger and more numerous going down the groups because the molecules are getting bigger with a greater number of electrons being dispersed. (1)

London dispersion forces: The weak forces of attraction between all atoms and molecules, caused by temporary dipoles.

Dipole Definition: A dipole is a separation of electrical charges.

EXTENDED RESPONSE• Q3) a) …Write an ion-electron equation for the first ionisation of potassium.

(1)

K(g) K+(g) + e’ (1)

Ionisation energy: The energy required to remove one mole of electrons from one mole of atoms in the gaseous state.

Remember the ionisation number lets you know the charge on the ion after electrons have been removed.

EXTENDED RESPONSE• 3) b) Explain Clearly why the first ionisation energy of potassium is smaller

than that of chlorine. (2)

• Distance from nucleus and valence electron increases and potassium has more electron shells than chlorine (i.e. potassium has an outer shell further from the nucleus), electrostatic interaction decreases as a result of distance. (1)

• Electron Shielding by core electrons will be greater for potassium compared to chlorine, valence electron requires less energy to remove from a potassium atom as a result. (1)

Ionisation energy: The energy required to remove one mole of electrons from one mole of atoms in the gaseous state.

Electron shielding: The shielding effect describes the attraction between an electron and the nucleus in any atom with more than one electron shell. Shielding effect can be defined as a reduction in the effective nuclear charge on the electron cloud, due to a difference in the attraction forces of the electrons on the nucleus.

EXTENDED RESPONSE• Q4) a) i) …name an element in the second period that exists as a covalent

network.

Carbon or boron (1)

Bonding in first twenty elementsyou must learn these!!!!

EXTENDED RESPONSE• Q4) a) ii)…why do the atoms decrease in size from lithium to neon? (1)

Number of occupied electron shells remains the same, in this case two but across the period nuclear charge increases and this greater force of attraction from the nucleus draws the electrons closer to the nucleus. This decreases the atomic size (1)

Covalent radius is half of the internuclear separation between the nuclei of two single-bonded atoms of the same species. (atomic size) to find bond length then add covalent radius values.

EXTENDED RESPONSE• Q4) a) iii) …Which element in the second period is the strongest reducing

agent.

Lithium (1)

Reducing agent: a substance that tends to bring about reduction by being oxidized and losing electrons. (low electronegativity).

Electronegativity: is a measure of the tendency of an atom to attract a bonding pair of electrons. The Pauling scale is the most commonly used. Fluorine (the most electronegative element) is assigned a value of 4.0, and values range down to caesium and francium which are the least electronegative at 0.7.

EXTENDED RESPONSE• Q4) b)On descending Group 1 from lithium to caesium, the electronegativity of the elements decreases. • Explain clearly why the electronegativity of elements decreases as you go down the group. (2)

• Electron shielding reduces nuclear attraction towards bonding electrons. (1)

• Distance is greater so attraction of nucleus is reduced when covalent bonding. (1)

Electronegativity: is a measure of the tendency of an atom to attract a bonding pair of electrons. The Pauling scale is the most commonly used. Fluorine (the most electronegative element) is assigned a value of 4.0, and values range down to caesium and francium which are the least electronegative at 0.7.

Electron shielding: The shielding effect describes the attraction between an electron and the nucleus in any atom with more than one electron shell. Shielding effect can be defined as a reduction in the effective nuclear charge on the electron cloud, due to a difference in the attraction forces of the electrons on the nucleus.

EXTENDED RESPONSE• Q5) Complete the table below by adding the name of an element for each of

the types of bonding and structure described. (2)

Use data book (melting points)

Bonding and structure at room temperature and pressure Name of element

Metallic solid sodium

Monatomic gas He, Ne, Ar, Kr, Xe, Rn

Covalent network solid C, B, Si

Discrete covalent molecular gas H2, N2, O2, F2, Cl2

Discrete covalent molecular solid P4, S8, I2, At2

EXTENDED RESPONSE• Q6) a) …Write an ion-electron equation for the first ionisation of sodium. (1)

Na(g) Na+(g) + e’ (1)

Ionisation energy: The energy required to remove one mole of electrons from one mole of atoms in the gaseous state.

Remember the ionisation number lets you know the charge on the ion after electrons have been removed.

EXTENDED RESPONSE• Q6) b) Explain clearly why the first ionisation energy of sodium is much

lower than its second ionisation energy. (3)The second ionisation in comparision to the first ionisation energy • will be the energy required to remove an electron

from an inner electron shell which is closer to the nucleus. (1)

• Removing from a positive ion instead of a neutral atom (1)

• Less electron shielding will be present when comparing the 2nd to the 1st ionisation, the first describing energy to remove the electron from the third shell. (1)

Ionisation energy: The energy required to remove one mole of electrons from one mole of atoms in the gaseous state.

Electron shielding: The shielding effect describes the attraction between an electron and the nucleus in any atom with more than one electron shell. Shielding effect can be defined as a reduction in the effective nuclear charge on the electron cloud, due to a difference in the attraction forces of the electrons on the nucleus.

EXTENDED RESPONSEQ7) The melting point of sulfur is much higher than that of phosphorous. Using your knowledge of chemistry, explain fully, in terms of the structures of sulphur and phosphorous molecules and the intermolecular forces between molecules of each element, why the melting point of sulfur is much higher than that of phosphorous. (3)

• Phosphorous has a melting point of 44°C and Sulphur has a melting point of 113°C• Both molecules exist as discrete covalent molecular solids at room temperature.

• Phosphorous has four atoms bonded with strong covalent bonds in a tetrahedral structure. • Sulfur can exist in different forms (allotropes) but the most common is 8 sulfur atoms joined in a ring by strong covalent

bonds. • Both have London dispersion forces between molecules but sulfur molecules are bigger with more electrons and

therefore will have stronger forces to overcome and this would explain it having a higher melting point that phosphorous.

EXTENDED RESPONSEQ8) The Periodic Table groups together elements with similar properties. In most Periodic Tables hydrogen is placed at the top of Group 1, but in some it is placed at the top of Group 7. Using your knowledge of chemistry, comment on the reasons for hydrogen being placed above either Group 1 or Group 7. (3)

Although the student has given two correct pieces of information regarding hydrogen, his/her answer contains no detail and therefore only shows a limited understanding. It is important to recognise that giving two pieces of information does not equate to two marks; these questions are designed to test a deep understanding of chemistry and this answer only gives the briefest response. The student gains 1 mark.

EXTENDED RESPONSEQ8) The Periodic Table groups together elements with similar properties. In most Periodic Tables hydrogen is placed at the top of Group 1, but in some it is placed at the top of Group 7. Using your knowledge of chemistry, comment on the reasons for hydrogen being placed above either Group 1 or Group 7. (3)

This answer contains a lot of correct information, and the student has shown a reasonable understanding of hydrogen as an element. However, the answer lacks a more in-depth discussion of the chemistry. The student gains 2 marks.

EXTENDED RESPONSEQ8) The Periodic Table groups together elements with similar properties. In most Periodic Tables hydrogen is placed at the top of Group 1, but in some it is placed at the top of Group 7. Using your knowledge of chemistry, comment on the reasons for hydrogen being placed above either Group 1 or Group 7. (3)

This answer provides ample evidence that the student has a good understanding of the chemistry of hydrogen. The answer includes information about electron arrangement, ion formation, bonding in molecular hydrogen and the reactivity of hydrogen gas, amongst other things. This student would gain 3 marks.