Different Elements Classifications

S.3 Classification of Elements

​Optional Extension: Documentary

​It’s good to be familiar with common elements on the periodic table. View the samples and explore the visual periodic table below to get an idea for what pure samples of each element can look like. Note that He, Ne, Ar, Kr and Xe are colorless gases and only light up if you pass electricity through them.

Periodic Table

The periodic table is arranged in increasing order of atomic number (number of protons).

​Groups: The vertical columns. The group number tells you how many valence electrons.

​Periods: Horizontal rows. The period number tells you how many energy levels it has.

​Group:

V. elec:

​Period 1

​Period 2

​Period 3

​Period 4

Tip: Group numbers tell you valence electrons.
Tip: Period numbers tell you number of energy levels.

​1

​2

​3

​4

​5

​6

​7

​8

​1

​2

​13

​14

​15

​16

​17

​18

​Note: The periodic table actually looks like this:

But to make it fit on paper, we usually cut out the actinides and lanthanides and stick them below:

​Metals, Non-metals and Metalloids

The elements are also grouped by properties:

1. Metals: Shiny, malleable, conduct electricity/heat. Solid at room temp, except mercury (Hg).

2. Non-metals: Dull, brittle, insulators. Mostly gases, one liquid (Br) and solid (C, P, S, I, Se).

3. Metalloids (semi-metals): Average properties between metals and non-metals.

Categorize

Options (11)

Li

Cu

Al

H

Xe

C

Si

B

Sb

Pb

S

Organize these options into the right categories

Metal

Non-metal

Metalloid

​Blocks

​Based on the electron configurations of atoms, the periodic table is also arranged into blocks according to sublevels.

Periodic Table – Royal Society of Chemistry

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Determining Electron Configuration

An alternative to the method learned earlier is to count along the periodic table if you memorize the blocks. Use whichever method you prefer.

1. ​Identify the element of interest in the periodic table.

2. Starting with hydrogen, move along each row, left to right, top to bottom writing out the electron configuration for each block in a period until you reach your element.

​Example: Titanium

Open Ended

Try using the block method to write the electron configuration of argon (Ar).

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​Periodicity

Trends in properties of elements across a period and down a group.

This is the big picture that we will try to understand over the next few slides. This image is printed later in your notebook, so you do not need to copy it down.

​Electrostatic Attraction

​All of chemistry happens because of the laws of physics governing attraction and repulsion between charged particles.

​Likes repel, so electrons (e-) repel each other.

​Opposites attract, so the negative electrons (e-) are electrostatically attracted to the positive protons (p+) in the nucleus.


Extension: Wondering why protons stick together in the nucleus even though they are positive and should repel? Search or ask AI about the strong nuclear force to find out!

​​Electrostatic Attraction in the Atom

Effect of Distance from the Nucleus

​Analogy: When opposite poles of magnets get further apart, their attraction weakens.

What it means: Electrons in energy levels further from the nucleus experience weaker attraction to the nucleus (and are therefore easier to remove).

Effect of Number of Protons (Nuclear Charge)

Analogy: The more magnets you use to attract a paperclip, the stronger you will attract it.

What it means: Electrons in the same energy level get more attracted to nuclei with more protons (and are therefore harder to remove).

IMPORTANT

• From the graphs above, you can see increased distance decreases attraction at a faster rate than increased protons increases attraction.

• This means when both distance and numbers of protons are changing, distance is the predominant factor.

Multiple Choice

Electrons are...

• negative and repel each other

• positive and repel each other

• negative and attract each other

• Positive and attract each other

Multiple Choice

Electrons are attracted to...

• other electrons

• protons

• neutrons

• none of the above

REVIEW: Ionisation Energy: X_(g) → X⁺_(g) + e^- _(g)

​The energy required to remove one electron from one atom in the gaseous state (→ ∞). HINT: If you forget the gaseous state, you lose a mark!

​In simpler terms: the harder it is to pull an electron away from its nucleus, the higher its IE is.

​Activity: Use the values for ionization energy from the data booklet and graph.
Skill: See if you can identify trends.
Skill: Identify discontinuities in trends.

​First Ionisation Energy

​

​Note the values for ionization energy can be obtained from the data booklet and graphed:

​First Ionisation Energy

​

As you go down a group, the ionization energy decreases.

Trend

​Explanation: As you go down a group, the number of energy levels increases, leading to increased distance between the valence electron and the nucleus. Since increased distance weakens the attraction between protons and electrons, the valence electron becomes easier to remove (lower IE).

​

Multiple Choice

Which element in Group 16 would have the lowest first ionisation energy?

• Oxygen

• Sulfur

As you go across a period, the ionisation energy increases.

Trend

​Explanation: As you go across a period, the nuclear charge (number of protons) increases while the number of energy levels remains constant. The increased nuclear charge increases the attraction to the valence electrons, making them harder to remove (higher IE).

​

Multiple Choice

Which element in Period 3 would have the highest first ionisation energy?

• Sodium

• Magnesium

• Aluminum

• Chlorine

• Argon

​Explanation: The valence electron for Be is in the 2s sublevel, while B’s valence electron is in the 2p sublevel. As the 2p sublevel is further from the nucleus than the 2s sublevel, B’s valence electron is easier to remove (lower IE). The same type of reasoning would apply to Al and Mg…

​

​B

​Be

Discontinuities:
Evidence for the existence of sublevels.

• B has lower 1st IE than Be

• Al has lower 1st IE than Mg

Open Ended

Why is the 1st ionisation energy of aluminum lower than magnesium?

• O has lower 1st IE than N

• S has lower 1st IE than P

​Explanation: The valence electrons in the 2p sublevel of N are all in half-filled orbitals. However, the last valence electron for O is in a fully filled 2p orbital (2 electrons). Since electrons are both negative and thus repel one another, the pairing in the same orbital makes this electron easier to remove (lower IE). The same type of reasoning would apply to S & P.

​

​N

​O

Discontinuities
Evidence for the existence of orbitals.

Open Ended

Why is the 1st ionisation energy of sulfur lower than phosphorus?

​Atomic Radius

​Distance from the nucleus to valence electrons of a neutral atom.

​Atomic
Radius

​Activity: We can also graph from the data booklet... What trends do you notice?

​

​Atomic Radius

​Explain why atomic radius increases as you go down a group:

As you go down a group, the number of energy levels increases. When there are more energy levels, the valence electrons are further away from the nucleus and therefore the atom gets larger.

​Atomic Radius

​Explain why atomic radius decreases as you go across a period:

As you go right across a period, the number of energy levels remains constant while the number of protons (nuclear charge) increases. The increase in nuclear charge increases the electrostatic attraction on the valence electrons, pulling them in closer (decreasing the atomic radius).

​Ionic Radius

Distance from the nucleus to valence electrons of an ion.

• The trends are the same as atomic radius except non-metals get larger.

​Ionic Radius

Usually questions will focus on comparing the radius of an ion to its neutral atom:

​METAL IONS (CATIONS)

Ionic radius < atomic radius

​Explanation: Metals lose their valence electrons, so their ions have one less energy level compared to their neutral atom. This makes them smaller in radius.

​Ionic Radius

​​NON-METAL IONS (ANIONS)

Ionic radius > atomic radius

​Explanation: Nonmetals gain more valence electrons to fill their valence shell. The extra electrons in the valence shell repel one another, causing them to spread out, making the ion larger compared to its neutral atom.

​Combining Ionic & Atomic Radii

​Atoms and ions that have the same electron configuration are called isoelectronic. We are sometimes asked to compare their radii.

​Explanation: The greater number of protons leads to a greater attraction to the valence electrons, pulling the electrons closer to the nucleus (which decreases the radius).

​

​10

​Electron Affinity

​The energy released when a neutral atom gains 1 electron to form a negative ion.

​Electron Affinity

​The energy released when a neutral atom gains 1 electron to form a negative ion.

​The stronger the force of attraction between the nucleus and the electron, the higher the electron affinity.

Values for first electron affinity are always negative because energy is released (exothermic). This will be covered more in a future unit.

​Electron Affinity

Period Trend: Increases as you go across a period because nuclear charge increases, causing a stronger attraction to electrons.

​Group Trend: Decreases as you go down a group because the number of energy levels increases, resulting in the electron being gained at an energy level that is further away from the nucleus and thus experiencing weaker attraction.

Drag and Drop

From highest electron affinity to lowest electron affinity: ​

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>​

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​ >​

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F

S

Cu

Sr

​Electronegativity

​How strongly an atom attracts a pair of bonding electrons in a covalent bond.

• ​The stronger the force of attraction between the nucleus and bonding electrons, the higher the electronegativity.

• Trends & explanations are the same as electron affinity.

​Summary of Trends

​NOTE: As you can have the data booklet for the entire exam, it's likely the new exams will focus on you explaining rather than just comparing.

​Metallic Character

​The more easily an atom loses an electron, the greater its metallic character. In other words, it is inversely proportional to an atom’s ionization energy (IE). The lower the IE, the greater the metallic character.

Open Ended

• Explain why the reactivity of metals increases down the group.

Drag and Drop

TREND: As you go down group 1, the reactivity (metallic character) of the alkali metals ​

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increases

decreases

remains constant

​Non-metallic Character

​The more easily an atom gains an electron, the greater its non-metallic character. It is proportional to its electron affinity. In other words, the higher the electron affinity’s magnitude, the higher the non-metallic character is.

​Non-metallic Character

Dropdown

As you go up group 7, the reactivity of the halogens ​​

Select answer

.

Open Ended

Explain why the reactivity of the group 7 elements increases as you go up the group.

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HL Material

​Transition Metals: Incomplete d-sublevels

​Transition metals have incomplete d-sublevels, which can help explain some of their properties.

​Variable Oxidation States

For example, iron forms two common ions:

​Variable Oxidation States Cont'd

​​Transition metals first lose their 4s electrons, but the 3d electrons are similar in ionization energy, so it is not uncommon to lose some of those at the same time. This leads to variable oxidation states.

High Melting Points

Transition metals have higher melting points because they have more delocalized electrons from the 4s and 3d sublevels, which leads to stronger metallic bonding. Metallic bonding will be explained in detail in the next unit.

​Magnetic Properties

Experiment: Magnets & different metals.

When electrons are paired & have opposite spins, the magnetic fields cancel out. Atoms with unpaired electrons in an orbital exhibit some magnetic properties. The more unpaired electrons an atom has, the more magnetic it is. Transition metals can have more unpaired electrons in the 3d sublevel, leading some to be more magnetic than other common elements.

​Catalytic Properties

​The 4s and 3d electrons in transition metals can sometimes form covalent bonds (see next unit in detail) with reactants. These weaken the bonds in the reactants, allowing the reaction to happen more quickly.

​Experiment: Cobalt(II) chloride reaction with H₂O₂

​Complex Ions & Colored Solutions/Compounds

Experiment: Observe samples of transition metal compounds/solutions.​
Ligands are molecules or ions in a solution that form coordination bonds with metal cations. This will be covered in more detail in the acids/bases unit.

​Formula: [Cu(H₂O)₆]²⁺

​Complex Ions & Colored Solutions/Compounds

​Ligands in solutions or anions in compounds can cause splitting of the d-sublevel in transition metals. The difference in energy between the split sublevels corresponds to energies of light in the visible spectrum.

​Complex Ions & Colored Solutions/Compounds

• ​Thus when an electron absorbs energy, that color is removed and you see the complementary color to what was removed.

• For example, this solution absorbs yellow. Light observed from the solution lacks yellow light, which our eyes & brain interpret as violet (the complementary color to yellow).

​Complex Ions & Colored Solutions/Compounds

• ​Thus when an electron absorbs energy, that color is removed and you see the complementary color to what was removed.

• For example, this solution absorbs yellow. Light observed from the solution lacks yellow light, which our eyes & brain interpret as violet (the complementary color to yellow).

​A transition metal complex appears purple and absorbed yellow light with a wavelength of approximately 𝜆 = 580 nm. Calculate the following, expressing your answers to 3 significant figures:

b) Frequency of light absorbed.

Special Note: Zinc is not a transition metal.

​The technical definition of a transition metal is that it forms ions with partially filled d-sublevels.
Zinc forms the Zn²⁺ ion, which has a full 3d sublevel.

​This results in it forming colorless solutions (due to no d-sublevel splitting).

​

​

​

Special Note: Sc²⁺ forms colored solutions, but Sc³⁺ forms colorless solutions.

Scandium is technically considered a transition metal because one of its ions forms colored solutions.

Multiple Choice

Which solution would be colorless.

Strategy: Look for transition metals = color.

• A

• B

• C

• D