Chem 14 Reviewer PDF

ATOMIC STRUCTURE Chemistry - A branch of science that deals with the study of the nature of the matter Matter - Anything that occupies space and has mass Molecule - The smallest particle that a compound can be reduced to before it breaks down into its element Element - A subs...

ATOMIC STRUCTURE Chemistry - A branch of science that deals with the study of the nature of the matter Matter - Anything that occupies space and has mass Molecule - The smallest particle that a compound can be reduced to before it breaks down into its element Element - A substance that cannot be decomposed any further by chemical action Compound - A combination of two or more elements Atom - Smallest form of a chemical particle that retains the properties of the particle - Comes from the Greek words "atomos" meaning "unable to be cut" - Original meaning was the smallest, invisible form of a chemical particle - An Atom is the smallest invisible particle of element, having all the characteristics of the parent element, which can neither be created nor destroyed by any chemical change. - It cannot exist freely. - It is the ultimate particle of an element, which may or may not have independent existence. - All elements are composed of atoms. - EARLY ATOMIC THEORIES The Indivisible Atom **Democritus** - Proposed that matter consisted of particles that could not be cut - Atoms are all composed of the same substance, all with different sizes and shapes, but with qualitative differences in weight, arrangement and position **Empedocles** - Affirmed that there are four fundamental permanent elements earth, fire, water & air - "The elements were held together by love & driven apart by strife" **Dalton Atomic Theory** **John Dalton (1766-1844)** - Posited the Atomic **Theory of Matter** which was based on the following 5 postulates.  - THREE LAWS OF MATTER **Law of Conservation of Mass** - Atoms Simply rearrange during chemical reaction - The atoms retain their masses - There is neither gain nor loss of mass during reaction - The mass of the substances before reaction is the same after the reaction - "NOTHING COMES FROM NOTHING" **Law of Definite Proportion** - If a compound is broken down into its constituent elements, the massed of the constituents will always have the same proportions, regardless of the quantity or source of the original substance **Law of Multiple Proportions** - If two elements can combine to form two or more compounds, the masses of one element that combines with a fixed mass of another are in the ratio of small whole numbers  - MODERN ATOMIC THEORIES **Sir Joseph John Thomson** - Proposed the plum pudding model of the atom. According to the model, it consists of a positively charged mass interspersed with electrons. **Ernest Rutherford** 1. was one of J.J. Thomson\'s students 2. proposed that the atom has nucleus, small and very dense and all of the mass of the atom was concentrated there. - FUNDAMENTAL PARTICLES OF AN ATOM **Proton** - The proton is a positively charged particle. - It has unit positive charge and unit mass. The mass of proton is approximately equal to the mass of one hydrogen atom. It is equal to 1.00732 amu. - The proton is present in atoms of all the elements. - The protons are present inside the nucleus of an atom. **Electron** - The electron is a negatively charged particle. - It has unit negative charge and negligible mass - The mass of an electron is about 1/1837 of the mass of a hydrogen atom. - Electrons are present in all the atoms. - Electrons are revolving around the nucleus in various circular orbits (shell) **Neutron** - The neutron is a neutral particle. Hence, it has no charge. - It has unit mass. The neutron is present in atoms of all elements except hydrogen. The mass of a neutron is slightly greater than the mass of a proton. It is equal to 1.00871 amu. - Neutron is present inside the nucleus of an atom **Atomic Number** - Based on the carbon standard the atomic mass of an element may be defined as the ratio between the mass of one atom of the element and 1/12th of mass of an atom of carbon - Atomic number= No of protons=No of electrons **Mass Number / Atomic Weight / Atomic Mass** - The mass number of an element is given by the total number of protons and neutrons present in the nucleus of an atom - Mass No. = (P+N) ; the number of neutrons is = Mass No. - Atomic Number **Isotopes** - The isotopes are atoms of the same elements having the same atomic number but different mass number **Isobar** - Isobars are the atoms of different elements having the same mass number but different atomic number LESSON 2 - PURE SUBSTANCE **Elements** - There are 118 known elements - Names of elements are sometimes derived from Greek or Latin words (bromine comes from the Greek word for stench), or the name of a place (e.g. the elements californium and germanium) or individual (e.g. einsteinium, fermium). - The symbol for an element is a chemist's shorthand which helps in writing chemical equations. - The symbol for sodium, Na, comes from the Latin natrium; the K for potassium comes from kalium; the Fe for iron is ferrum. **Metal** - Almost 80% of the elements are metals. Characteristic: ***Luster*** - Metals reflect light and are shiny ***Conductivity**.* They are good conductors of heat and electricity. ***Ductility***. They can be drawn into thin wires. ***Malleability***. They can be flattened into thin sheets. **Non-Metal** - Is characterized by the absence of metallic properties. - Is characterized by the absence of metallic properties. - Can be gas, liquid, or solid. - They aren't shiny (lustrous) - They don't conduct heat or electricity well. **Metalloids** - Elements that exhibit some properties of metal and non-metals under specific conditions. - Have also the ability to conduct electricity, but not as well as metals. - Are brittle and easily broken **Compound** - Is a substance that is composed of a fixed proportion of two or more elements. - It can be broken down (or decomposed) into its constituent elements by chemical techniques. **Organic Compound** - Are those which are formed by carbon and hydrogen or carbon and hydrogen with oxygen, nitrogen, and a few other elements. - Carbon atoms are very important because they are found in living organisms. - Can be obtained from natural sources, such as plants and animals. **Inorganic Compound** - Is a substance that does not contain both carbon and hydrogen - Derived from nonliving components, and generally have ionic bonds, lack carbon hydrogen bonds, and rarely, if ever, contain any carbon atoms. **Acids** - Are substances that are generally sour in taste and are present when blue litmus paper turns red and neutralizes bases. **Bases** - Substances that can neutralize acids are called bases. - It has a bitter taste and is present when red litmus paper turns blue. **Salts** - Is a neutral substance whose aqueous solution does not affect litmus. - Substances formed when acids react with bases.  **LESSON 3** - **MOLECULES & IONS** **Molecule** - Is the smallest particle of matter (element or a compound) that can exist freely. - The molecule is made up of two or more atoms of the same element or different elements. It can be further divided into atoms. **Homoatomic Molecules** - The molecule is made up of two or more atoms of the same elements. - Most of the elementary gases consist of homoatomic molecules. (hydrogen, oxygen) **Heteroatomic Molecules** - The molecule is made up of two or more atoms of different elements. **Ions** - **Any atom or group of atoms that bears one or more positive or negative electrical charges.** - **Cations -- positively charged Ions Anions -- negatively charged Ions** **CHEMICAL FORMULA** **Molecular Formula** - Is the short form of representation (symbolical) of one molecule of an element or a compound. **Significance of molecular formula** 1\. It shows the elements present in one molecule. 2\. It gives the exact number of atoms present in one molecule. 3\. It is used to calculate the molecular mass of a molecule. **Empirical Formula** - Is the relative ratios of different atoms in a compound. - Simplest form of molecular formula **Structural Formula** - representation of how atoms are bonded to one another in a molecule CHEMICAL BONDING **Chemical Bonds** - Is a lasting attraction between atoms, ions or molecules that enables the formation of chemical compounds. **Metallic Bonds** - Exists among metals such as iron, gold, silver, magnesium, zinc or copper. - Metals are composed of a lattice of atoms of the same element where the electrons are loosely held and are freely moving or "delocalized" - Are strongest of all bonds between atoms. **Octet Rule** - An atom will be most stable when surrounded by 8 electrons in the valence shell - Atoms lose, gain or share electrons in order to have a full valence shell of 8 electrons **Ionic Bonds** - Also known as Electrovalent Bond - Is formed as a result of the complete transfer of one or more electrons from one atom to another  **Covalent Bonds** - Is formed by the mutual sharing of pair of electrons between two atoms with each atom supplying the equal number of electrons for sharing.  **Molecular Mass** - Molecular mass of an element or a compound is the ratio between the mass of one molecule of the element or the compound and the mass of ½ part of a carbon atom. **Lesson 3** **Dobereiner's Triads** -**German Chemist Johann Wolfgang Dobereiner** group the elements into three which he called "triads". Law of Dobereiner's Triads - Some groups of three elements ('triads') shared common physical properties. - The mean or average of the atomic weights of the heaviest and lightest elements was roughly or exactly equal to the atomic weight of the middle element. - Some other physical properties (e.g.. density) also showed this trend. The average of the relevant property of the heaviest and lightest elements was very close to that of the middle element. Examples of Dobereiner's Triads The molecule is made up of two or more atoms of the same elements. Most of the elementary gases consist of homoatomic molecules. (hydrogen, oxygen) Alkaline Earth Triad  Alkali Metal Triad Halogen Triad  Chalcogen Triad **Newlands' Law of Octaves** In 1865, John A. Newlands (british chemist) proposed a new classification of elements. Newlands found that if elements were arranged to increasing atomic weight, chemical properties repeated in an interval of eight. Newlands Octaves arranged the known 62 elements then into roughly 8 groups of 8 each. The elements were in increasing order of atomic weight. Every element shared properties with the 8th element after it. Sometimes two elements were made to occupy one place, to reflect their properties.  **Limitations of Newlands' Law of Octaves** It left no space for new elements to be discovered. It grouped some dissimilar elements together, e.g. halogens with metals. It worked well only up to Calcium Sometimes two elements had to take up one place, with no explanation. **Mendeleev's Periodic Table** **Dmitri Mendeleev (1834-1907)** "Elements arranged according to the size of their atomic weights show clear periodic properties." Predicted Gallium, which in his time he called it eka-aluminum by observing its density, and atomic weight of 68. Also predicted scandium, germanium, rhenium. **Moseley's Periodic Table and the Periodic Law** **Henrey Mosley (1913)** "When the elements are arranged according to increasing atomic numbers, certain properties repeat periodically." Experiments using x-ray helped him discover atomic numbers which make him saw that the errors in Mendeleev's table could be corrected. Atomic Mass of cobalt is slightly large than that of nickel.  Lesson 4 GASES Gases behave differently from one another. However, they also manifest common properties. Properties of gases may be grouped into two -- the general properties and measurable properties. General Properties of Gases ❑ Most gases exist as molecules, usually diatomic. ❑ Gases have no definite shape and volume. ❑ Gases are easily compressed when pressure is applied. ❑ Gases expand when heated and contract when cooled. ❑ Gases exert pressure. ❑ Densities of gases are relatively small compared to the densities of solids and liquids. Measurable Properties of Gases **Pressure** Pressure is defined as force per unit area. The pressure of a gas is the force exerted by the gas on the walls of its container divided by the surface area of the container. Mathematically, pressure is expressed as The standard unit of pressure under System International (SI) is Pascal (Pa) which is equivalent to a force of one newton ( **1 N = 1 kg m/s2**) acting on an area of one square meter. **1atm = 760 torr = 760mmHg** **1atm = 1.013 bar = 14.7 psi** **1atm = 101 325 Pa or 101.325 kPa** Where: atm = atmosphere, mmHg = millimeter mercry Pressure is measured using an instrument called barometer, developed by Evangelista Torricelli, an Italian Physicist. **Volume** The volume of a gas is the space it occupies. Also, the volume of a vessel is equal to the volume of the gas it contains. This is based on the principle that gas particles occupy all the spaces available. The common units of volume used in gas measurements are: **cubic meter (m3)** **cubic centimetre (cm3)** **Liter (L)** **Milliliter (mL)** **Temperature** Temperature of a gas is usually expressed in three units, the degree Celsius (oC), degree Fahrenheit (oF), and kelvin (K). **Only Kelvin is used in computations involving** **temperature of a gas.** **oC = 5/9 (oF -- 32)** **oF = 9/5 (oC +32)** **K = oC +273** **Amount of Gases** The quantity of the gas being measured is always expressed in moles or n. Units of mass such as kilogram and gram should be converted to moles. The formula needed to do this is  Density The density of a gas computed by dividing the mass of the gas by its volume. The unit often used is gram per liter (g/L) **Standard Conditions, STP** The temperature of a gas and atmospheric pressure vary from place to place and from time to time and since the volume of a gas is dependent on its temperature and pressure, it is important to have a set of standard conditions for these quantities. This set of standard conditions is called STP or standard temperature and pressure. **Standard pressure = 1 atm or 760 torr** **Standard temperature = 00C or 273K** **One mole of a gas occupies a volume of 22.4L** **Gas Laws** **Boyle's Law** Named after the british chemist, Robert Boyle, this is the gas law that relates pressure and volume. It states that at constant temperature, the volume of a gas is inversely proportional to the pressure applied. **PV = k** **P1V1 = P2V2** Where: P1= initial pressure P2= final pressure V1 = initial volume V2 = final volume Charles' Law Jacques Charles, states that the volume of gas and the applied absolute temperature is directly proportional at constant pressure. **Avogadro's Law** Named after the Italian Scientist, Lorenzo Amadeo Avogadro, states that the volume of a gas and number of moles present is directly proportional at constant temperature and pressure. (n = is the number of moles of gas) 𝑉 = kT **Combined Gas Law** ⬥ When Boyle's law and Charle's Law are considered together, the resulting principle is Combined Gas Law. It state that "for a given mass of gas, the volume is inversely proportional **Ideal Gas Law** Combining the gas laws will result in a single equation. This is the equation that relates all the four variables (P, V, n and T). All gases whose behaviour is approximated by this equation are called ideal gases. **PV= nRT** Where R is the proportionality constant called the gas constant. The value of R derived from standard temperature and pressure (00C and 1 atm) is **0.0821L∙atm / mol ∙ K or** **8.3145 m3 ∙ Pa / mol ∙k or** **8.3145 J / mol ∙k**