SCH3U1 Unit 1 Study Notes PDF

Unit 1: Summary of Early Theories of Matter: Democritus: - Believed atoms were uniform, solid, hard, incompressible, and indestructible and that they moved in infinite numbers through empty space until stopped Aristotle: - Believed all materials on Earth were not made...

Unit 1: Summary of Early Theories of Matter: Democritus: - Believed atoms were uniform, solid, hard, incompressible, and indestructible and that they moved in infinite numbers through empty space until stopped Aristotle: - Believed all materials on Earth were not made of atoms, but of the four elements, Earth, Fire, Water, and Air - All substances were made of small amounts of these four elements of matter Dalton: - Atomic Theory - All matter is composed of extremely small particles called atoms. - Atoms of a given element are identical in size, mass, and other properties - Atoms of different elements differ in size, mass, and other properties. - Atoms cannot be subdivided, created, or destroyed. Thomson: - Believes all atoms contain tiny negatively charged subatomic particles or electrons. Nagaoka: - Proposed a model of the atom that contained a small nucleus surrounded by a ring of electrons Rutherford: - The positive charge and most of the mass of an atom is concentrated in an extremely small volume Chadwick:  - Atoms consisted not only of protons and electrons but also neutrons - Shot particles through gold Subatomic Particle Relative Mass Charge Location Proton 1 1- Nucleus Neutron 1836.12 1+ Nucleus Electron 1838.65 0 Orbitals Terms: Atomic Number (Z): Numbers of protons in atomic numbers Mass Number (A): Sum of protons and neutrons Isotopes: Atoms that have the same number of protons but different neutrons Relative Atomic Mass (Aᵣ): The weighted of atoms (in grams) of an element Unified Atomic Mass Unit (u): A unit of mass used to express atomic masses Atomic number - Atomic Mass = Neutrons Atomic radius: The measurement from an atom’s nucleus to the furthest orbital of electrons. Why does atomic radii decrease from left to right across each period? Valence electrons are being added to the same energy level at the same time the nucleus is increasing in protons. Why does the atomic radius increase from top to bottom in a given group? Down a group, the number of energy levels (n) increases, so there is a greater distance between the nucleus and the outermost orbital. What two particles are responsible for most of the mass of an atom? Neutrons and protons. What is the shielding effect? The decrease in the nucleus's force of attraction on valence electrons due to the existence of electrons in the inner shells. Why is it more prominent in elements at the bottom of the group? The nuclear core is farther removed from the valence electrons and the atomic radius increases as it goes down. Define ionic radius. The distance from the nucleus of an ion up to which it has an influence on its electron cloud. When an atom loses an electron to become a positive ion (cation), its radius will decrease. Explain why. It will lose valence electrons causing the radius to shrink. When an atom gains an electron to become a negative ion (anion), its radius will increase. Explain why. There will be a greater electron reputson Why is an Al3+ ion smaller than a Na+ ion? Sodium has less electrons causing there to have a less powerful inward force on the electrons. Define ionization energy. Be sure to include the difference between first ionization energy and second ionization energy. The amount of energy required to remove an electron from an isolated atom or molecule. First ionization energy: The energy needed to remove the outermost, or highest energy, electron from a neutral atom in the gas phase. Second ionization energy: The energy required to remove the outermost, or least bound, electron from a 1+ ion of the element What is the trend for ionization energy as you move across a period? Why? Ionization energy becomes greater due to the size of the ions across the period getting smaller. What is the trend for ionization energy as you move down a group? Why? Ionization energy decreases as you move down due to atomic size becoming larger as you go down. Define electron affinity. The change in energy (in kJ/mol) of a neutral atom (in the gaseous phase) when an electron is added to the atom to form a negative ion Explain why some electron affinity values are positive and some are negative. A negative electron affinity indicates that the process absorbs energy, increasing the energy of the system A positive electron affinity indicates that the process releases energy, lowering the energy of the system. What is the trend for electron affinity as you move across a period? Why? Electron affinity increases meaning it becomes more negative the more right you go due to greater nuclear attraction. What is the trend for electron affinity as you move down a group? Why? Electron affinity decreases meaning it will become more negative due to the electrons being placed in a higher energy level far from the nucleus, thus a decrease from its pull. Define electronegativity. A chemical property that describes the tendency of an atom or a functional group to attract electrons toward itself What is the trend for electronegativity as you move across a period? Why? Electronegativity increases as you move across a period due to nuclear change becoming larger. What is the trend for electronegativity as you move down a group? Why? Electronegativity decreases as you move down a group due to atomic number increasing. Properties of Ionic and Covalent Compounds Type of Compound Property IONIC MOLECULAR Type of Elements Present metal + non-metal non-metal + non-metal Type of Bonding ionic / covalent ionic / covalent Electron Behaviour in Bond transferred / shared transferred / shared Smallest Unit of Compound formula unit / molecule formula unit / molecule State at Room Temperature solid / liquid / gas solid / liquid / gas Melting Point relatively high / low relatively high / low Electrical Conductivity yes / no yes / no Solubility in Water relatively high / low relatively high / low  electrolyte Electrolyte is the mineral inside your body that are able to conduct electricity. intramolecular force Intramolecular force is an attractive force that brings atoms and ions together in a compound. intermolecular force Intermolecular force is an attractive force within molecules. What properties of ionic compounds suggest that ionic bonds are strong? - High melting and boiling points - Structure and tend to be rigid and brittle Predict the charge on the most stable ion formed by each of the following elements. Indicate the ion by writing the symbol complete with charge. (a) sulfur -2 (b) barium +2 (c) bromine -1 (d) chlorine -1 (e) calcium +2 (f) potassium +1 (g) phosphorus -3 diatomic molecule A diatomic molecule are two atoms that have the same element or different elements in which a bonded chemically together. polyatomic molecule A polyatomic molecule is 3 or more atoms that are together within a firm shape. lone pair A lone pair are two valence electrons which don’t share with other atoms in a covalent bond. octet rule The octet rule is the likelihood of atoms to want to have 8 valence electrons in its shell. structural formula The structural formula helps to know where the chemical bonds are within a molecule and its atoms.  bonding capacity The bonding capacity is how many chemical bonds there can be in an atom within a chemical element. Ionic and covalent: ΔEN = 0 (non-polar covalent bond) ΔEN < 1.7 (polar covalent bond) ΔEN > 1.7 (ionic bond) van der Waals Forces London dispersion force dipole-dipole force Hydrogen bonding The London dispersion The dipole-dipole force is The hydrogen bonding is force is the least powerful an attractive force that is in with a unique type of intermolecular force. the middle of two polar dipole-dipole force in the Making a transient. It is a molecules, one at the end middle of molecules. This short term attractive force and one at the negative occurs by an attractive and occurs when two atoms end. force within a hydrogen close together create atoms atom that is covalently that make temporary bonded to an atom that is dipoles. extremely electronegative. - The strongest intermolecular force that can exist between covalent molecules is dipole-dipole force. - The weakest intermolecular force between covalent molecules is Londot dispersion force. All types of molecules participate in this type of bonding. When a strong intermolecular force is present between covalent molecules, this will result in higher melting and boiling points because it takes more energy to break the intermolecular forces of attraction between molecules. When a weak intermolecular force is present between covalent molecules, the substance will likely be gas and the molecules are not very attracted to each other. How to find boiling point: 1. Find electronegativity 2. Subtract second term by first Higher electronegativity = higher boiling point Example: CaBr₂ Electronegativity of Ca= 1 Electronegativity of Br₂= 2.96 2.96 - 1 = 1.96