Chapter 3: The Periodic Table Flashcards
Explain how elements in the periodic table are arranged
1) By increasing atomic number
2) By period (horizontal rows)
3) By group (vertical columns)
4) By metal/metalloid/non-metal
- Metal: Alkali metals, metals, transition metals
- Metalloids: have properties of both metals and non-metals (demarcated by ‘staircase’ boundary at B)
- Non-metals: halogens and noble gases
State the properties of a metal
Metallic structure: lattice of cations in a sea of delocalised electrons
- Metals: able to lose electrons easily (low ionisation energies) to form cations
- Across a period: ionisation energies increase, elements lose electrons less readily (metals usually formed by elements on left side of periodic table)
- Down a group: ionisation energies decrease, elements lose electrons more readily (metals usually formed by elements at the end of a group (eg. group 14)
METALS:
1) Large atomic radii
2) Low ionisation energies
3) Low (exothermic) electron affinity
4) Low electronegativity
Use position of element on periodic table to determine its electron configuration
(Eg. Sulfur: Period 3, Group 16)
- Highest filled energy level: 3
- Number of electrons at highest filled energy level: 6
3s²3p⁴
Explain trend in atomic radii
(down the group and across a period)
ATOMIC RADIUS: size of an atom (half the internuclear distance of a molecule of that element)
Determined by:
1) Nuclear charge: number of protons (creating “attractive force”)
2) Shielding effect: repulsion of valence electrons by inner electrons (number of “core” non-valence electrons)
3) Effective nuclear charge (nuclear charge - shielding effect)
Increases down a group:
- Atoms have more electron shells down a group
- Nuclear charge and shielding effect both increase, counteracting any effect caused by a greater number of protons
Decreases across a period:
- Nuclear charge increases (no. of protons increase)
- Shielding effect remains the same (no. of non-valence electrons remain the same)
- Effective nuclear charge increases (outer electrons are more strongly attracted to the nucleus)
- Atomic radii decreases
Explain trend in ionic radii (cationic and anionic radii relative to their atomic radii), and across an isoelectronic series
IONIC RADIUS: size of an ion
Cationic radii: smaller than its atomic radii
- Cations have fewer electron shells than their atoms
- Fewer electrons
- Less repulsion between electrons, valence electrons are more strongly attracted to the nucleus
- Smaller electron cloud
Anionic radii: larger than its atomic radii
- Constant nuclear charge
- More electrons
- More repulsion between electrons
- Larger electron cloud
Across isoelectronic series (Si⁴⁻, P³⁻, S²⁻, Cl⁻):
- Increase in nuclear charge
- Constant screening effect
- Increase in effective nuclear charge
- Valence electrons are more strongly attracted to nucleus
- Ionic radii decreases
Explain trends in first ionisation energy
(down the group and across a period)
FIRST IONISATION ENERGY: energy required to remove the outermost electron from a gaseous atom
M (g) –> M⁺ (g) + e⁻
Decreases down the group:
- Increase in atomic radius
- Valence electrons are less strongly attracted to the nucleus
- Less energy required to remove the outermost electron
Increases across a period:
- Increase in nuclear charge
- Constant screening effect
- Increase in effective nuclear charge
- Smaller atomic radius
- Valence electrons are more strongly attracted to the nucleus
- More energy required to remove the outermost electron
*Exception across Period 2:
1) Group 2 (Be) and Group 13 (B)
- B has a smaller AR, and is supposed to have higher IE. However, because the electron being removed from B is at the 2s sub-shell and the electron removed from Be is at the 2p sub-shell (higher energy level), B has a lower IE than Be
2) Group 15 (N) and Group 16 (O)
- O has a smaller AR and is supposed to have higher IE. However, because O has 2 paired electrons in the 2p-orbital generating interelectronic repulsion whereas N only has 1 electron
(2p³), O has a lower IE
Explain trends in electron affinity
(down the group and across a period)
ELECTRON AFFINITY: energy change when one electron is added to a gaseous atom
X (g) + e⁻ –> X⁻ (g)
- Usually exothermic for all elements: bond forming is exothermic
Decreases down a group:
- Increase in AR
- Electron will be less strongly attracted
- Releases less energy when electron is brought in
- Less exothermic
*Exception: Group 17
Cl has the most exothermic electron affinity (supposed to be F)
- However, due to the small size of the F atom, it generates large amounts of interelectronic repulsion, requiring more energy to attract an electron and releasing less energy
Increase across a period:
- Decrease in AR
- Electron will be more strongly attracted to nucleus when brought in
- Releases more energy
- More exothermic
*Exception: Period 3
- P (group 15) is supposed to have more exothermic electron affinity than Si (group 14)
- However, P has 3 electrons in 3 separate p-orbitals (3p³), which means that bringing in an extra electron would generate interelectronic repulsion, weakening the attraction between the nucleus and the electrons, making its electron affinity less exothermic
Explain trends in electronegativity
(down a group and across a period)
ELECTRONEGATIVITY: measure of the attraction of an atom in a molecule for the electron pair in the covalent bond of which it is a part
Decreases down a group:
- Increase in AR
- Increase in bonding distance between nucleus and electron pair
- Weaker attraction between electron pair and nucleus
- Decreased electronegativity
Increases across a period
- Decrease in AR
- Decrease in bonding distance between nucleus and electron pair
- Stronger attraction between nucleus and electron pair
- Increased electronegativity
State the properties and reactions of Group 1 metals (alkali metals)
All elements down a group:
- Similar chemical properties
- Gradation in physical properties
Group 1 metals:
- Metallic bonding: electrostatic attraction between the lattice of cations and the sea of delocalised electrons
Melting point decreases down the group:
- IR increases
- Nucleus is further away from delocalised electrons
- Weaker attractive force between nucleus and delocalised electrons
- Less energy required to break apart the lattice and overcome the IMFA between particles
REACTIONS:
- Alkali metals are all highly reactive:
1) With oxygen (to form an oxide)
2) With water (to form a base and hydrogen gas)
- Get more reactive down the group: AR increases, IE decreases
State the properties and reactions of Group 17 elements (halogens)
- Diatomic molecules
- Bonding in halogens: covalent bonding (weak intermolecular forces of id-id) to form simple molecular structure
Boiling point increases down the group:
- Mᵣ increases
- London forces between molecules grow stronger
- More energy required to overcome them
Reactivity decreases down the group:
1) With alkali metals (to form salts)
2) With other halogens (displacement reaction)
State the different types of oxides and describe the changes in types of oxides across a period
OXIDES:
- Basic (react with water to form alkaline solutions IF SOLUBLE + react with acids to form salts)
- Acidic (react with water to form acidic solutions IF SOLUBLE + react with bases to form salts)
- Amphoteric (react with both acids and bases to form salts + does not react with water)
- Neutral
Across a period:
Basic (metallic oxides) –> Amphoteric (some metal oxides) –> Neutral (some non-metallic oxides) -> Acidic (non-metallic oxides)
Write the equations for the reactions of oxides with water and predict the acidity of the resulting solutions
Sodium oxide (Na₂O): Basic
Na₂O + H₂O –> 2 NaOH
Magnesium oxide (MgO): Basic
MgO + H₂O –> Mg(OH)₂
Aluminium oxide (Al₂O₃): Amphoteric
Al₂O₃ + 6H⁺ –> 2Al³⁺ + 3H₂O
Al₂O₃ + 2OH⁻ + 3H₂O –> 2Al(OH)₄⁻
Phosphorous pentoxide (P₄O₁₀): Acidic
P₄O₁₀ + 6 H₂O –> 4 H₃PO₄
Sulfur dioxide (SO₂)/sulfur trioxide (SO₃):
SO₂ + H₂O –> H₂SO₃
SO₃ + H₂O –> H₂SO₄
