Applied Science Unit 1 Chemistry Knowledge Organiser

Electronic Orbitals

Electronic orbitals refer to regions in an atom where there is a high probability of finding electrons. These orbitals are defined by quantum numbers and are integral to understanding electron configuration in atoms. The types of orbitals include:

• s orbitals: Spherical in shape; can hold a maximum of 2 electrons.
• p orbitals: Dumbbell-shaped; can hold a maximum of 6 electrons.
• d orbitals: More complex shapes; can hold a maximum of 10 electrons.
• f orbitals: Even more complex shapes; can hold a maximum of 14 electrons.

Aufbau Principle

The Aufbau principle is a guideline for determining the electron configuration of atoms. According to this principle, electrons fill orbitals starting from the lowest energy level to the highest. This means:

1. Electrons will first occupy the 1s orbital,
2. Then fill the 2s orbital,
3. After that, they will fill the 2p orbitals, and so on. This continuous filling follows the order determined by their increasing energy levels.

Bonding Types

Ionic Bonding

Ionic bonding occurs when electrons are transferred from one atom to another, resulting in the formation of ions. This typically happens between metals and non-metals where:

• Metals lose electrons to become positively charged ions (cations).
• Non-metals gain electrons to become negatively charged ions (anions).
  The resulting electrostatic attraction between the oppositely charged ions forms an ionic bond. An example is the formation of sodium chloride (NaCl).

Covalent Bonding

Covalent bonding involves the sharing of electrons between atoms to achieve a full outer shell. This typically occurs between non-metal atoms. The shared electrons allow each atom to attain a noble gas configuration, resulting in the formation of molecules such as water (H₂O) and carbon dioxide (CO₂).

Metallic Bonding

Metallic bonding is characterized by a 'sea of electrons' where electrons are delocalized and shared among a lattice of metal cations. This explains the properties of metals:

• Electrical conductivity due to the mobility of the delocalized electrons.
• Malleability and ductility, enabling metal to be shaped without breaking.

Bohr Theory

The Bohr theory describes the structure of the hydrogen atom, proposing that:

• Electrons travel in fixed orbits around the nucleus without radiating energy.
• Energy levels in these orbits are quantized.
• An electron moves to a higher orbit by absorbing energy and falls to a lower orbit by emitting energy, which corresponds to specific wavelengths of light (spectral lines).

Intermolecular Forces

Intermolecular forces, or forces between molecules, are responsible for many physical properties of substances. Key types include:

• London dispersion forces: Weak forces that arise due to temporary dipoles in molecules.
• Dipole-dipole interactions: Occur between molecules with permanent dipoles.
• Hydrogen bonds: A strong type of dipole-dipole interaction involving hydrogen bonded to highly electronegative atoms (N, O, F).

Empirical Formula

The empirical formula of a compound gives the simplest whole-number ratio of atoms of each element present. For example, for hydrogen peroxide (H₂O₂), the empirical formula is HO, indicating a 1:1 ratio of H to O.

Reacting Quantities

Reacting quantities refer to the amounts of reactants and products involved in a chemical reaction. This is typically expressed using chemical equations and stoichiometry, where the coefficients represent the ratio of reactants to products.

Periods 1, 2, 3, and 4

In the periodic table, periods are horizontal rows that signify the number of electron shells in the atoms of the elements found in those rows. For example:

• Period 1 (H, He): 1 shell
• Period 2 (Li, Be, B, C, N, O, F, Ne): 2 shells
• Period 3 (Na, Mg, Al, Si, P, S, Cl, Ar): 3 shells
• Period 4 (K, Ca, Sc, Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn): 4 shells

Groups: s Block, p Block, d Block

s Block

The s block comprises groups 1 and 2 of the periodic table, which includes alkali metals and alkaline earth metals. These elements have their outermost electrons in s orbitals.

p Block

The p block contains groups 13 to 18 and includes a diverse range of elements, including nonmetals, metalloids, and metals. Their outermost electrons are in p orbitals.

d Block

The d block corresponds to transition metals in groups 3 to 12. These elements are characterized by the filling of d orbitals and exhibit properties like variable oxidation states and a wide range of oxidation states.

Percentage Yield

Percentage yield is a useful metric in chemistry that describes the efficiency of a reaction. It is calculated using the formula:
$ext{Percentage Yield} = \frac{ ext{Actual Yield}}{ ext{Theoretical Yield}} imes 100 \%$
This measurement indicates how much product was obtained from a reaction relative to how much was theoretically possible based on stoichiometry.

Molecular Formula

The molecular formula of a compound indicates the actual number of atoms of each element present in a molecule. For example, the molecular formula of glucose is C₆H₁₂O₆, demonstrating it contains 6 carbon, 12 hydrogen, and 6 oxygen atoms.

Atomic Radius

The atomic radius is a measure of the size of an atom, typically defined as the distance from the nucleus to the outermost electron shell. The atomic radius tends to increase down a group and decrease across a period due to increased nuclear charge and electron shielding.

Electronegativity

Electronegativity is the measure of an atom's ability to attract shared electrons in a chemical bond. The scale of electronegativity is defined by Pauling, with values typically ranging from 0.7 (Fr) to 4.0 (F). An increase in electronegativity results in stronger polar covalent bonds.

First Ionization Energy

First ionization energy is the energy required to remove the outermost electron from a neutral atom in its gaseous state. This energy tends to increase across a period (due to increasing nuclear charge) and decrease down a group (due to electron shielding).

Electron Affinity

Electron affinity is the energy change that occurs when an electron is added to a neutral atom in the gaseous state. A more negative electron affinity indicates a greater tendency for an atom to gain an electron.

Ionic Radius

The ionic radius refers to the size of an ion. Cations (positively charged ions) are generally smaller than their parent atoms due to loss of electrons, while anions (negatively charged ions) are larger due to the addition of electrons and increased electron-electron repulsion.

Type of Bonding

The type of bonding in compounds can be classified as ionic, covalent, or metallic depending on the elements involved and their electronegativities. Understanding the type of bond helps predict the physical and chemical properties of the substances.

Melting and Boiling Points

The melting and boiling points of substances are influenced by the strength of intermolecular forces. Generally, stronger forces (like ionic and hydrogen bonds) lead to higher melting and boiling points, while weaker forces (like London dispersion forces) result in lower melting and boiling points.

Reactivity Series

The reactivity series is an order of metals and nonmetals based on their reactivity. It predicts which metals will displace others in chemical reactions. A common series from most reactive to least reactive is:

• Potassium (K)
• Sodium (Na)
• Calcium (Ca)
• Magnesium (Mg)
• Aluminium (Al)
• Carbon (C)
• Zinc (Zn)
• Iron (Fe)
• Lead (Pb)
• Copper (Cu)
• Silver (Ag)
• Gold (Au)

Products and Reactivity of Period 2 and 3 Elements with Oxygen

The elements in periods 2 and 3 react with oxygen to form oxides, which can be either ionic or covalent depending on the element involved. For instance:

• Period 2 Example: Lithium (Li) reacts with oxygen to form lithium oxide (Li₂O).
• Period 3 Example: Phosphorus (P) can form phosphorus pentoxide (P₂O₅).

Variable Oxidation States

Variable oxidation states refer to the ability of certain elements, primarily transition metals, to exist in more than one oxidation state. This property is significant in redox reactions and coordination compounds.

Displacement Reactions

Displacement reactions involve a more reactive element displacing a less reactive element from its compound. An example would be:
Zn + CuSO_4
ightarrow ZnSO_4 + Cu
Here, zinc displaces copper from copper(II) sulfate due to its higher reactivity.

Uses and Applications of Substances

Substances derived from chemical reactions have numerous applications ranging from industrial processes to household products. For example:

• Ammonia (NH₃) is produced from nitrogen and hydrogen and is used in fertilizers.
• Sodium chloride (NaCl), common salt, is used in food preservation.

Oxidation and Reduction

Oxidation refers to the loss of electrons or an increase in oxidation state, while reduction refers to the gain of electrons or a decrease in oxidation state. These processes are fundamental to redox reactions where oxidation and reduction occur simultaneously. A mnemonic to remember this is LEO says GER (Lose Electrons is Oxidation, Gain Electrons is Reduction).