What's the equation for (effective) nuclear charge?

atomic number - outer electrons = effective nuclear charge

outer electrons - atomic number = effective nuclear charge

What does this image show?

The periodic table is arranged into 4 blocks (s,p,f,d)

The periodic table consists of groups and periods

The number of valence electrons can be deduced from the group that an element is in

The period number(n) is the outer energy level occupied by electrons

What does this image show?

Shows the positions of metals, non-metals and metalloids

The periodic table is arranged into 4 blocks (s,p,f,d)

The periodic table consists of groups and periods

The effective nuclear charge...

Increases across the period, stays the same down the group

Decreases across the period, stays the same down the group

Increases across the period, decreases down the group

Decreases across the period, increases down the group

As you move across a period, the atomic and ionic radius...

Only the atomic radius increases and the ionic radius decreases

Only the ionic radius increases and the atomic radius decreases

As you move down a group, the atomic and ionic radius...

Only the atomic radius increases and the ionic radius decreases

Only the ionic radius increases and the atomic radius decreases

Which is correct for ionization energy?

It increases across a period and decreases down a group

It decreases across a period and increases down a group

It increases across a period and stays the same down a group

It stays the same across a period and stays the same down a group

Why do the reactions of alkali metals with water become more vigorous further down the group?

Less energy is required to remove valence electrons

There's lower ionization energy further down the group

There's higher ionization energy further down the group

More energy is required to remove valence electrons

Why do reactions with the reactivity of halogens decrease down the group?

Because the atomic radius increases down the group making it less easy to gain an electron

Because the ionic radius increases down the group

Because less energy is required to remove valence electrons

Because ionization energy decreases down the group

Which is correct for electron affinity?

It increases across the period; more exothermic as needs less electrons to be stable

It decreases down the group; less exothermic as number of main energy levels increases

The atom decreases in size with the increase of core charge across the period

What is the exception for electron affinity? Why?

Carbon is more exothermic than nitrogen because of the half filled orbitals are more stable in nitrogen

Chlorine is more exothermic than fluorine

Chlorine is more endothermic than fluorine

Nitrogen is more exothermic than carbon because of the half filled orbitals are more stable in carbon

Which is correct for electronegativity?

It increases across the period; the core charge increases and the atomic radius decreases

It decreases down the group; number of main energy levels and shielding effect increases

The atom decreases in size with the increase of core charge across the period

What are the similarities and differences in properties of elements in the same group?

All alkali metals LOSE electrons because they're in group 1

Halogens are all ACIDIC and generally have the HIGHEST electronegativity in their period

All halogens LOSE one electron because they're in group 7

Alkali metals are all ACIDIC and generally have the HIGHEST electronegativity in their period

a group of very reactive metals in group 1 of the periodic table

a group of reactive non-metals in group 7 of the periodic table

a group of very non-reactive gases in group 0 of the periodic table

a group of metals in the d-block of the periodic table

a group of reactive non-metals in group 7 of the periodic table

a group of very non-reactive gases in group 0 of the periodic table

a group of very reactive metals in group 1 of the periodic table

a group of metals in the d-block of the periodic table

a group of very non-reactive gases in group 0 of the periodic table

a group of reactive non-metals in group 7 of the periodic table

a group of very reactive metals in group 1 of the periodic table

a group of metals in the d-block of the periodic table

Define transition metals

a group of metals in the d-block of the periodic table

a group of reactive non-metals in group 7 of the periodic table

a group of very non-reactive gases in group 0 of the periodic table

a group of very reactive metals in group 1 of the periodic table

a series of 14 elements in the periodic table; generally ranging in proton number from cerium to lutetium

elements 90 to 103, they are unstable and undergo radioactive decay

a column of the periodic table which contains elements with similar chemical properties; same no. electrons in their outer shell

a horizontal row in the periodic table which contains elements with same number of shells, and with an increasing no. of electrons in the outer shell

Define actinides/actinoids

elements 90 to 103, they are unstable and undergo radioactive decay

a column of the periodic table which contains elements with similar chemical properties; same no. electrons in their outer shell

a horizontal row in the periodic table which contains elements with same number of shells, and with an increasing no. of electrons in the outer shell

a series of 14 elements in the periodic table; generally ranging in proton number from cerium to lutetium

Re-write this definition of group: a column of the periodic table which contains elements with similar chemical properties; same number of electrons in their outer shell

a column of the periodic table which contains elements with similar chemical properties; same number of electrons in their outer shell

Re-write this definition of period: a horizontal row in the periodic table which contains elements with the same number of shells, and with an increasing number of electrons in the outer shell

a horizontal row in the periodic table which contains elements with the same number of shells, and with an increasing number of electrons in the outer shell

Re-write this definition of ligand: a molecule or negative ion that donates a pair(or pairs) of electrons to a central metal ion to form a dative/coordinate covalent bond

a molecule or negative ion that donates a pair(or pairs) of electrons to a central metal ion to form a dative/coordinate covalent bond

Re-write this definition of electronegativity: a measure of the tendency of an atom in a molecule to attract a pair of shared electrons towards itself

a measure of the tendency of an atom in a molecule to attract a pair of shared electrons towards itself

which is NOT a chemical property of transition metals

have variable oxidation states as their valence electrons are in two different sets of orbitals

act as a catalyst when either as an element or compound

Which is NOT a physical property of transition metals

high electrical and thermal capacity

Why isn't zinc considered a transition metal?

It doesn't form ions with incomplete d-orbitals

What does this image show?

transition elements show an oxidation state of +2 when the s-electrons are removed

one of the arrow from 4s moves to the end of 3d

transition elements show an oxidation state of +1 when the s-electrons are removed

Why is scandium not considered a transition metal

The ion(Sc3+) doesn't have an incomplete d-sub level but the atom does

It has variable oxidation states

Which is FALSE for complex ions?

The type of bond formed between metals and ligands are ionic bonds

Ligands donate a pair of electrons to a central metal ion with a specific charge

Creates a charged particles of a central metal atom with either 2,4 or 6 ligands

Total charge = charges of the ligands + charges of the central metal ion

Which is TRUE for magnetic properties of transition metals?

The more unpaired electrons, the stronger the magnetic property (Paramagnetic)

Unpaired electrons exist when the complex has an odd number of electrons or because electron pairing is destabilized

The more paired electrons, the weaker the magnetic property (Diamagnetism)

Why do transition metals appear colored? Re-write: Visible light is absorbed as the d-orbitals split into 2 sub levels. Electrons get excited from low to high energy levels, thus complementary color is reflected

Visible light is absorbed as the d-orbitals split into 2 sub levels. Electrons get excited from low to high energy levels, thus complementary color is reflected

What can affect of colors of transition metal ion complexes?

Nature of the ligand: strong field ligands causes a large split in d-orbital energies of the central metal ion

Oxidation state of the metal: as it oxidation state increases, so does the amount of splitting of the d-orbitals, therefore it changes the color of light absorbed

Geometry of the complex: usually 6 coordinate complexes are less intensely colored than 4 coordinate complexes

What can be identified from the spectrochemical series?

The ligands are arranged in order of their ability to split d-orbitals in transition metal complexes

The greater the splitting, the LESS energy difference between the two sets of d-orbitals so LESS energy absorbed

It goes from strong field ligands to weak field ligands

I- has the shortest wavelength and CO has the longest wavelength

Which is the imposter?

Wavelength of light absorbed increases if the stronger field ligand is replaced with a weaker one

Red has the highest wavelength and violet has the lowest wavelength

Wavelength of light absorbed decreases due to the large energy difference between the two sets of d-orbitals

When the color of light is absorbed, the complimentary color is transmitted/ reflected

IB Chemistry HL - Periodicity

Quiz

•

Chemistry

•

11th - 12th Grade

•

Practice Problem

•

Medium

•

NGSS

HS-PS1-1, HS-PS1-7, HS-PS4-3

+4

Standards-aligned

QuackerzUwU QuackerzUwU

Used 32+ times

FREE Resource

50 questions

Show all answers

1.

MULTIPLE CHOICE QUESTION

30 sec • 1 pt

Which is correct about Na2O?

It's a simple strongly acidic oxide

The oxide ion, O2-, is a very weak base so it has a low tendency to combine with H ions

It's a simple strongly basic oxide

The reaction of Na2O with H2O gives a pH of 7

2.

FILL IN THE BLANK QUESTION

1 min • 1 pt

Na2O(s) + H2O(l) --> ?

Tags

NGSS.HS-PS1-7

3.

MULTIPLE CHOICE QUESTION

30 sec • 1 pt

Which is correct for MgO?

The reaction of MgO with HCl gives a pH of around 10.3

It requires less energy to break

It's more strongly basic than Na2O since oxide ions are more weakly-bound

The reaction of MgO with water gives a pH of around 10.3

Tags

NGSS.HS-PS1-1

4.

FILL IN THE BLANK QUESTION

1 min • 1 pt

MgO(s) + H2O(l) --> ?

Tags

NGSS.HS-PS1-7

5.

MULTIPLE SELECT QUESTION

30 sec • 1 pt

Which is correct for P4O10

It is classified as a basic oxide

It has a pH of 0-1

It reacts violently with water to give a very weak acid of H2PO4 (phosphoric acid)

It has a pH of 13-14

6.

FILL IN THE BLANK QUESTION

2 mins • 1 pt

Construct the balanced equation:
NO2 reacts with water to form nitrous acid(HNO2) and nitric acid(HNO3)

7.

FILL IN THE BLANK QUESTION

2 mins • 1 pt

Construct the balanced equation:
NO2 reacts with water but the nitrous acid(HNO2) decomposes to nitric oxide(NO) with nitric acid(HNO3)

Tags

NGSS.HS-PS1-1

NGSS.HS-PS1-2

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IB Chemistry HL - Periodicity