Chemistry Patterns (OCR)

1. The Periodic Table

Development of the Periodic Table:

• The periodic table was developed by Dmitri Mendeleev in 1869. He arranged the elements based on atomic mass and observed that elements with similar properties occurred at regular intervals. He also left gaps in his table for undiscovered elements, predicting their properties.

• Henry Moseley (1913) later rearranged the periodic table according to atomic number instead of atomic mass, resolving some inconsistencies in Mendeleev's table.

Periodic Law:

• The periodic law states that the properties of elements recur periodically when they are arranged by increasing atomic number. As you move across periods or down groups in the periodic table, there are observable trends in properties like atomic radius, ionization energy, and electronegativity.

Groups and Periods:

• Groups: Vertical columns in the periodic table, with elements in the same group sharing similar chemical properties because they have the same number of valence electrons. For example, Group 1 (alkali metals) all have one valence electron.

• Periods: Horizontal rows of the periodic table, which represent the filling of electron shells. Properties change across a period due to the increasing nuclear charge.

Periodic Trends:

• As you move left to right across a period, atomic size decreases, ionization energy increases, and electronegativity increases.

• As you move down a group, atomic size increases, ionization energy decreases, and electronegativity decreases.

2. Atomic Structure and Electron Configuration

Subatomic Particles:

• Protons have a positive charge (+1) and are located in the nucleus of the atom.

• Neutrons have no charge (neutral) and are also found in the nucleus.

• Electrons have a negative charge (-1) and are found in electron shells outside the nucleus.

Electron Configuration:

• Electrons are arranged in energy levels or shells around the nucleus. The first shell can hold up to 2 electrons, the second up to 8, and so on. The distribution of electrons in the shells is called the electron configuration.

• Example: For carbon (C), which has an atomic number of 6, the electron configuration is:

  

  This shows that carbon has 2 electrons in the first shell and 4 electrons in the second shell.

Effective Nuclear Charge (Z_eff):

• The effective nuclear charge is the net positive charge felt by an electron in an atom. It takes into account the attraction from the protons in the nucleus and the repulsion from other electrons.

• Formula: 

Where:

• Z is the atomic number (number of protons).

• S is the shielding constant (which accounts for the shielding effect of inner electrons).

3. Trends in the Periodic Table

Atomic Radius:

• Across a Period: As you move from left to right across a period, the atomic radius decreases. This is because the number of protons (and hence the nuclear charge) increases, which pulls the electrons closer to the nucleus, reducing the size of the atom.

• Down a Group: As you move down a group, the atomic radius increases. This happens because new electron shells are added, making the atom larger, and the outermost electrons are further from the nucleus and experience more shielding from inner electrons.

Ionization Energy:

• Ionization energy is the energy required to remove one mole of electrons from one mole of gaseous atoms. This energy increases as you move across a period because the atoms become smaller and the electrons are held more tightly by the increased nuclear charge.

• Across a Period: Ionization energy increases due to the increased nuclear charge and smaller atomic radius.

• Down a Group: Ionization energy decreases because the outer electrons are further from the nucleus, and there is greater shielding from inner electrons.

  • Formula for First Ionization Energy: 

Electronegativity:

• Electronegativity is a measure of an atom's ability to attract and bond with electrons. It generally increases across a period and decreases down a group.

  • Across a Period: Electronegativity increases because atoms are smaller and the nuclear charge is stronger, so they attract electrons more strongly.

  • Down a Group: Electronegativity decreases because atoms are larger and the outer electrons are further from the nucleus, leading to weaker attraction for bonding electrons.

4. Chemical Bonding and Properties

Ionic Bonding:

• Ionic bonds form when one atom transfers an electron to another, typically between a metal and a nonmetal. The metal becomes a positively charged ion (cation), and the non-metal becomes a negatively charged ion (anion).

  • Example: In sodium chloride (NaCl), sodium (Na) loses one electron to become:

    

    and chlorine (Cl) gains that electron to become

    

    These oppositely charged ions are held together by electrostatic forces of attraction.

• Ionic Bonding Properties:

  • High melting/boiling points due to strong ionic bonds.

  • Conduct electricity in molten or dissolved form because the ions are free to move.

  • Brittle structure due to the arrangement of ions in a lattice, which can break easily under stress.

Covalent Bonding:

• Covalent bonds are formed when two nonmetal atoms share one or more pairs of electrons to achieve a full outer shell of electrons.

  • Example: In water (H₂O), each hydrogen shares one electron with oxygen, forming covalent bonds.

• Covalent Bonding Properties:

  • Low melting/boiling points compared to ionic compounds.

  • Poor electrical conductivity because they do not have free-moving charged particles.

  • Solubility in non-polar solvents (like benzene).

Metallic Bonding:

• In metallic bonding, metal atoms release their outer electrons to form a "sea of electrons" that are free to move throughout the metal structure. The positive metal ions are held together by the attraction to the delocalized electrons.

• Metallic Bonding Properties:

  • Good conductors of electricity because the delocalized electrons can move freely.

  • Malleable and ductile: Metals can be hammered into thin sheets or drawn into wires because the layers of metal ions can slide over each other.

  • High melting/boiling points due to strong metallic bonds.

5. Transition Metals

Characteristics:

• Transition metals have partially filled d-orbitals. They are typically harder, stronger, and more conductive than the main group metals.

• Multiple Oxidation States: Transition metals can lose different numbers of electrons, leading to multiple oxidation states. For example, iron can exist as

  

  and

  

Ligands and Coordination Chemistry:

• Transition metals form coordination complexes with ligands. A ligand is a molecule or ion that donates a pair of electrons to the metal ion. These complexes often result in different colors and reactivities.

  • Example: In the complex

    

    water molecules act as ligands, coordinating with the copper ion.

6. Reactivity Series

Reactivity Trends:

• The reactivity series arranges metals by their reactivity with water, acids, and other substances. Highly reactive metals, such as potassium (K), are at the top, and less reactive metals, such as gold (Au), are at the bottom.

• Displacement Reactions:

  • More reactive metals displace less reactive metals from their compounds. 

  • For example:  

    

  • Here, zinc displaces copper from copper sulfate, forming zinc sulfate and copper metal.

7. Group Trends

Group 1: Alkali Metals:

• Alkali metals (e.g., lithium, sodium, potassium) react vigorously with water to form hydroxides and hydrogen gas. The reactivity increases down the group because the outer electron is more easily lost as the atomic radius increases.

  • General Reaction with Water: 

Group 7: Halogens:

• Halogens (e.g., chlorine, bromine, iodine) are highly reactive non-metals that tend to form salts with metals. Their reactivity decreases down the group because the atomic radius increases and the ability to gain an electron decreases.

  • Displacement Reaction: 

Group 0: Noble Gases:

• Noble gases (e.g., helium, neon, argon) are inert because they have a full outer electron shell, making them chemically stable and unlikely to form compounds.