Notes on Inorganic Chemistry PDF

Chemistry lies near the heart of many matters of public concern: Lesson 1: SCIENCE ADVANTAGES Is the study of and accumulated and accepted knowle...

Chemistry lies near the heart of many matters of public concern: Lesson 1: SCIENCE ADVANTAGES Is the study of and accumulated and accepted knowledge, which has been systematized and formulated to serve as the basis for the ❖ Improvement of health care discovery of general truths. ❖ Conservation of natural resources Derived from the Latin word “SCIRE” which means to know or ❖ Protection of the environment knowledge. ❖ Provision of our everyday needs for food ❖ Clothing and shelter ❖ Discovered pharmaceutical chemicals that enhance our health and DIVISION OF SCIENCE prolong our lives. ❖ Increased food production through the development of fertilizers and 1. Social Science - deals with the study of human behavior. pesticides. ❖ Developed plastics and other materials that are used in almost every facet ❖ Anthropology – the study of man and his culture of our lives. ❖ Economics – the study of how man allocates his resources to meet his needs ❖ Sociology – the study of the origin and constitution of society. DISADVANTAGES 2. Natural Science – deals with the study of natural objects and ❖ Some chemicals also have the potential of harming our health in the phenomena. This branch of science includes Physical and environment. Biological Sciences A. Physical Sciences – deal with the predictable behavior of the world CLASSIFICATION OF CHEMISTRY around us and include: ❖ Astronomy – the study of heavenly bodies INORGANIC Inorganic Chemistry- deals with the general study of ❖ Chemistry – the study of the structure, properties and all the elements and their compounds, except those of composition of substances, and the changes that substances carbon. undergo. ❖ Geology – the study of earth ORGANIC deal with the study of carbon and its compounds. ❖ Mathematics – the study which deals with the relation of quantities. BIOCHEMISTRY specializes in the chemistry of the substances comprising living organisms. B. Biological Sciences – deal with the study of living things, their structure, processes, and influencing factors. Among them are: PHYSICAL the study of the structure of matter, energy changes, and the laws, principles and theories that govern the ❖ Microbiology – the study of microorganisms transformation of matter. ❖ Botany – the study of plants ❖ Chemistry – includes Organic and Biochemistry INDUSTRIAL deals with the conversion or transformation of raw ❖ Medicine – deals with the detection, prevention and cure of materials into finished products. diseases. ❖ Zoology – the study of animals ANALYTICAL concerns the identification, separation and quantitative determination of the composition of different substances. CHEMISTRY DIVISION OF ANALYTICAL CHEMISTRY It is the science that deals with the study of matter, its properties, and the changes in composition, in which matter undergoes the energy involved and the conditions necessary for the QUALITATIVE deals with the identification and composition of transformation of matter. substances. Determines WHAT elements or compounds are present in a given sample. AIMS OF CHEMISTRY QUANTITATIVE deals with the determination of the percentages or proportions these elements are present in the given ❖ To study the different properties of matter so as to distinguish one sample. This branch of Chemistry finds out HOW kind of matter from another. MUCH of these elements is present in a sample. ❖ To devise different methods of separating one substance from another such as silver from mercury in a solution, sugar from OTHER FIELD OF CHEMISTRY sugar cane, iron from iron ore, etc. ❖ To discover conditions necessary for promoting or preventing AGRICULTURE In Agriculture, Chemistry is utilized in soil analysis changes in matter such as the use of preservatives to prevent and in the manufacture of fertilizers. decomposition of food and the use of catalysts to influence change in matter. INDUSTRY new and better products are being produced from ❖ To devise different methods of converting natural raw products raw materials. into finished commodities necessary for the enjoyment of human life such as alcohol from sugar, paper from pulp, drugs and CHEMICAL ENG. a combination of chemistry and engineering that medicines from plants. Etc. improves or develops industrial processes for ❖ To determine the amount of malfunction of energy involved in a making commercial amounts of desirable chemicals given change such as the amount of energy furnished by a given that have been produced only in small quantities in quantity of steam. the laboratory.  COLLOID the study of the behavior of particles of matter that COMPOUND A pure substance made up of two or more CHEMISTYRY are larger than ordinary molecules but smaller than elements in a fixed ratio by mass. objects that can be seen with the best optical Formula of a compound tells us the microscope. ratios of its constituent elements and identifies each element by its atomic ELECTROCHEM the study of the chemical reactions that are symbol. produced by or that produce an electric current. NaCl: the ratio of sodium atoms to chlorine atoms in sodium chloride is NUCLEAR CHEM the study of radioactivity, the atomic nucleus and 1:1 nuclear reactions and the development of H2O: the ratio of hydrogen atoms to applications for radioactive isotopes in medicine oxygen atoms in water is 2:1 and industry. MATTER CLASSIFICATION OF MATTER MIXTURE A combination of two or more pure substances The substances may be present in any mass ratio. Each substance has a different set of physical properties. Mixtures may be homogeneous or heterogeneous. If we know the physical properties of the individual components of the mixture, we can use appropriate physical means to separate the mixture into its component parts. CHEMISTRY Chemistry is the study of matter. 1. INTENSIVE PROPERTY – is a property of matter that is – Matter is anything that has mass and takes up space. independent of the quantity of the substance. Matter can change from one form to another. Example: density, specific gravity – In a chemical change (chemical reaction), substances are used up and others formed in their place 2. EXTENSIVE PROPERTY – depends on the quantity of a Example: when propane (bottled or LP gas) burns in air, propane and oxygen are substance. converted to carbon dioxide and water. Example: mass, volume – In a physical change, matter does not lose its identity. – A common physical change is a change of state. Example: Ice (solid water) melts to become liquid water; liquid water boils to become steam (gaseous water). CLASSIFICATION OF MATTER ELEMENTS A substance (for example, carbon, hydrogen, and iron) that consists of identical atoms. – There are 118 known elements. – Of these, 88 occur in nature; the others have been made by chemists and physicists. – Their symbols consist of one or two letters. – Names are derived from a variety of sources: the English name of the element, people important in atomic science, geographic locations, planets, mythological sources, etc.  Lesson 2: Atom Modern Atomic Theories and Atomic Models is the smallest unit of matter that can retain the properties of an element. 1886 – E. Goldstein discovered the existence of protons as the positively It is made up of the positively charged particle (proton) and neutrally charged subatomic particle. 1897 – Sir JJ Thomson discovered that there are negatively charged charged particle (neutron) located at the nucleus and the negatively subatomic particles and called it the electron. Upon the discovery of the charged particle (electron) which is moving around the nucleus in a electron, Thomson formulated the first atomic model which is the quantized manner. “Plum Pudding Model” Early Atomic Theories MODERN ATOMIC THEORIES AND ATOMIC MODELS 5th Century BC Democritus of Abdera (5th Century) conceived the existence of an ultimate particle and called it “atomos” which means 1911 - Ernest Rutherford observed that the mass of an indivisible. atom is concentrated at the center. From this observation, he suggested the existence of a nucleus and formulated “laughing philosopher” the nuclear model of the atom. 4th Century BC – Plato conceptualized the existence of elements. 3rd Century BC – Aristotle developed the theories made by Plato and stated that “all substances were made up of matter on which different Sir JJ Thomson’s Plum Pudding Model of an Atom forms could be impressed. Modern Atomic Theories and Atomic Models ❖ All matter is composed of very tiny particles, which Dalton called Rutherford model of an atom atoms. ❖ All atoms of the same element have the same chemical properties. Atoms of different elements have different chemical properties. ❖ Compounds are formed by the chemical combination of two or more 1913 – Neils Bohr formulated the planetary model of an of the same or different kinds of atoms. atom. He stated that electrons found outside the nucleus ❖ A molecule is a tightly bound combination of two or more atoms that travel around the nucleus in definite orbits. acts as a single unit. Evidence for Dalton’s Atomic Theory Bohr’s Planetary Model of an Atom The great French Chemistry Antoine laurent Lavoisier (1743-1794) discovered the Law of Conservation of Mass, which states that matter can be neither created nor destroyed. Lavoisier proved this law by conducting many experiments in which he showed that the total mass of matter at the end of the experiment was the same as that at the beginning. Another French Chemist, Joseph Proust (1754-1826), demonstrated MODERN ATOMIC THEORIES AND ATOMIC MODELS the Law of Constant Composition, which states that any compound is always made up of elements in the same proportion by mass. Wave Mechanical Model Physicists namely Werner Heisenberg, Louis de Broglie and Erwin Schrodinger developed the wave-mechanical model of an atom which was based on the quantum mechanics developed by Max Planck and Albert Einstein ❖ 1924 – Louis de Broglie originated the idea that an electron has dual Evidence for Dalton’s Theory properties, that of a particle and a wave. ❖ Schrodinger – provided mathematical calculation of energy of each Law of Conservation of Mass electron in an atom. This model shows that electrons occupy various ❖ Matter can be neither created or destroyed. energy levels and sublevels. ❖ As Dalton explained, if matter is made up of indestructible atoms, then any chemical reaction just changes the attachments among atoms but Werner Heisenberg’s Uncertainty Principle – states that it is impossible to does not destroy the atoms themselves. simultaneously determine the exact position and the exact momentum of a body as ❖ Monatomic elements: consist of single atoms; for example, helium (He) small as an electron. and neon (Ne). 1932 – James Chadwick discovered the neutrally charged subatomic particle ❖ Diatomic elements: there are seven elements that occur as diatomic neutron which is also located at the nucleus of the atom. He also noted that the molecules: – H2, N2, O2, F2, Cl2, Br2, and I2 number of neutrons is equal to the number of protons in the nucleus. ❖ Polyatomic elements: some elements have three or more atoms per molecule: – O3, P4, S8 – diamond has millions of carbon atoms bonded together to form one gigantic cluster. Atomic Notation is a system of writing an atom showing its atomic number, atomic symbol and mass number Atomic symbol is the shorthand notation of an element. Usually, atomic symbols are letters taken from the Latin name of the elements Examples H – hydrogen O – oxygen  The main energy level is also alternately called the principal quantum number, orbits Atomic number is the number of protons present in an atom. Since an or electron shells (n). atom is neutral, the number of protons is equal to the Numbers 1, 2, 3, 4, 5, 6 and 7 were used to denote the distance of the electron to the number of electrons. nucleus. Examples n = 1 is the electron shell nearest to the nucleus and is the first atomic orbit Hydrogen has 1 proton therefore the atomic n = 2 is the second electron shell or the second atomic orbit number of hydrogen is 1 n = 7 is the farthest atomic orbit Nitrogen has 7 protons therefore the atomic number of nitrogen is 7. Each energy level can accommodate different number of electrons Mass Number is the sum of the number of protons and neutrons in an element. Examples: Helium has two protons and two neutrons, therefore, the mass number of helium is 4. Carbon has six protons and six neutrons, therefore, the mass number of carbon is 12. Within each principal energy level, electrons occupy energy sublevels called atomic orbitals or electron subshells. Different energy levels have different numbers of orbitals. Properties and Location within Atoms of Protons, Neutrons and Electrons The number of electrons each orbital can accommodate. Orbitals of an atom can be s, p, d or f. Different orbitals can accommodate different maximum numbers of electrons. The unit of mass is the atomic mass unit (amu) – By definition, 1 amu 1/12 the mass of a carbon atom with 6 protons and 6 neutrons. 1 amu = 1.6605 x 10-24 g The Aufbau Principle In distributing the electrons, the lowest energy levels are filled to the capacity before going to the next level. Hund's rule Electrons must fill up orbitals of the same energy level before pairing with another electron. ISOTOPES In some instances, the number of neutrons of an atom may be altered by radioactive means. This produces isotopes of an element. Isotopes – are atoms that have the same number of protons but different numbers of neutrons resulting in differences in the mass number of the same elements. Pauli’s Exclusion Principle States that no more than two can occupy a single ELECTRONIC CONFIGURATION orbital. ORBITAL – is the region in space around the nucleus of an atom is a designation of how electrons are distributed among various where the probability of finding the electron is greatest. orbitals in principal shells and subshells. to the fixed energy of electrons thus the term energy level. In order for electrons to move from one energy level to another Valence Shell and Valence Electrons energy level, it must gain or lose a fixed amount of energy. That fixed amount of energy to move an electron from one energy Valence – in an atom means the outermost Valence shell – is the outermost shell of an atom. It is the orbital that has the highest energy level to another is called quantum energy. level. Valence electrons – are the electrons found at the outermost shell of an atom.  Electron Dot or Lewis Structure of Atoms Five Methods of Representing Electron Distribution 1. Spectronic or spdf notation – the expected distribution of electrons using the spectronic notation indicates the energy levels and sublevels are filled. 2. Half –shell method – involves representing the distribution of electrons among energy levels. 3. Orbital diagram – the method is used to show how electrons are distributed among the orbital. 4. Electron dot formula / Lewis dot formula – the symbol of the element is written first and the valence electrons are represented by dots. 5. Noble gas notation – shorthand way using the inert gasses or noble gasses. Quantum Numbers Electrons of an element may be described by four quantum numbers: Principal Quantum number (n) – describes the number of electron clouds or electron shells/orbits. It is described as numbers 1-7. Secondary or azimuthal quantum number (l) – describes the shape of the orbital. Azimuthal quantum numbers may range from 0 to n-1. Electrons in an s orbital have an l value of 0, in the p orbital 1, in the d orbital 2 and in the f orbital 3. Magnetic quantum number (m) –describes the orientation of the electrons in the orbital. Values of m can range from –l, 0 or +l. If n = 1, the only possible value of m is 0. Magnetic spin quantum number (s) – describes the spin of the electrons. Values of s can only be +1/2 or 1/2 which describes clockwise and counterclockwise spin respectively. IONS are outcomes of either gain or loss of electrons loss of electrons in an atom produces a positively charged substance called cation. Gain of electrons in an atom produces a negatively charged substance called anion. Cations and anions are produced because of the behavior of atoms to gain or give off electrons to have 8 valence electrons called octet rule.  Lesson 3: PERIODICITY PERIODIC CLASSIFICATION OF THE ELEMENTS ❖ The long form periodic table or periodic chart is a list of all the known elements (118) arranged in order of increasing atomic number in horizontal rows of such a length that elements which are chemically alike fall directly beneath one another. ❖ It serves as a guide in predicting the electronic arrangement of atoms. The physical and chemical behavior of the elements can also be related to their position in the periodic table. DEVELOPMENT OF THE PERIODIC TABLE NEWLAND’S TABLE ( 1864) – John Newland developed a periodic table in CLASSIFICATION OF ELEMENTS which the elements were arranged in order of increasing atomic weights. ❖ He arranged the elements at intervals of eight, similar to the octave of REPRESENTATIVE ELEMENTS the musical scale. The term representative element is related to the stepwise addition of electrons to the s and p sublevels of the atoms. DMITRI MENDELEEV (1834-1907) Found in A families ❖ arranged the known elements in order of increasing atomic weight They exhibit a wide range of chemical behavior and physical beginning with hydrogen. characteristics ❖ He observed that when elements are arranged in this manner, certain Some of the elements are diamagnetic and some are paramagnetic sets of properties recur periodically. ❖ He then arranged elements with recurring sets of properties in the TRANSITION ELEMENTS same column (vertical row); Li, Na, and K, for example, fall in the same The columns 1B to 8B contain the transition elements or the B column and start new periods (horizontal rows). family/group. Group starts with 3B up to 8B which has 8 columns and then ends with 1B and 2B. These sequences which contain 10 elements each, are related to the stepwise addition of the 10 electrons to the d sublevel of the atoms. These elements are all metallic, dense, lustrous, good conductors of heat and electricity and in most cases, hard. INNER-TRANSITION ELEMENTS The additional horizontal rows comprise two groups of elements which were discovered to have similar characteristics as Lanthanum in the 6th period called Lanthanides (rare-earth elements) and Actinium in the 7th period called Actinides (heavy rare elements). These two series are called MEYER’S TABLE (1870) INNER-TRANSITION ELEMENTS. ❖ unaware of Mendeleev’s study, Julius Meyer had been working on his periodic table which consisted of 56 elements. He also maintained that NOBLE GASSES the properties of the elements were periodic functions of their atomic Colorless monatomic gasses, which are chemically unreactive and mass. diamagnetic (with the exception of the helium, which has the MOSELEY’S TABLE (1914) configurations of 1s2) ❖ The early periodic tables were arranged according to increasing atomic All the noble gasses have outer configurations of ns2np6, a very weights and this misplaced several elements, such as Ar, and K, Co and stable arrangement Ni, and Te and I, in the periodic table. Group 8 or group zero in old, group 18 in new. PERIODIC LAW DEVELOPMENT OF THE PERIODIC TABLE ❖ States that the physical and chemical properties of the elements are a John Dalton’s atomic periodic function of their atomic number. ❖ theory paved the way to the establishment of the periodic table. Dalton assigned weights and combining capacities to the atoms that by the early 19th century, the approximate atomic weights for more than 20 elements had been determined. DOBEREINER’S TRIADS (1829) – ❖ Johann Wolfgang Dobereiner grouped elements, which exhibit very similar characteristics in three or triads. The atomic weight of the second element was found to be the average of the first and third elements.  PERIODS, GROUPS AND FAMILIES The seven horizontal rows in the periodic table are called periods. ❖ Period 1 has 2 elements corresponding to 2 electrons in the s sublevel. ❖ Period 2 and 3 have 8 elements corresponding to 8 electrons in the s and p sublevels. ❖ Period 4 and 5 have 18 elements corresponding to 18 electrons in the s, p and d sublevels. ❖ Period 6 and 7 has 32 elements corresponding to 32 electrons in the s, p, d, and f sublevels. ❖ The vertical columns are called groups or families, which are divided into A and B subgroups. Some of the A families are designated by names. Group 1A – Alkali Metals Group 2A – Alkaline Earth Metals Group 7A – Halogens Group 8A – Noble Gases ❖ The other subgroups are classified according to the first element in the column. Group 3A – Boron Family Group 4A – Carbon Family Group 5A – Nitrogen Family Group 6A – Oxygen Family CLASSIFICATION OF ELEMENTS Metals ❖ are solids at room temperature (except for Hg, which is a liquid), shiny, Atomic Size conduct electricity, and are ductile and malleable. ❖ form alloys (solutions of one metal dissolved in another); brass, for ❖ The size of an atom is determined by the size of its outermost occupied example, is an alloy of copper and zinc. orbital. ❖ In chemical reactions, they tend to give up electrons. Example: the size of a chlorine atom is determined by the size of its Nonmetals three 3p orbitals, the size of a carbon atom is determined by the size of ❖ Except for hydrogen (H), they lie on the right side of the Periodic Table. if its three 2p orbitals. ❖ Except for graphite, do not conduct electricity. ❖ In chemical reactions, they tend to accept electrons. Metalloids ❖ They have some of the properties of metals and some of nonmetals; for example, they are shiny like metals but do not conduct electricity. ❖ Eight elements are classified as metalloids: boron, silicon, germanium, arsenic, antimony, polonium, astatine and tellurium. ❖ One of the metalloids, silicon, is a semiconductor; it does not conduct electricity under certain applied voltages, but becomes a conductor at higher applied voltages. Metallic Property ❖ The lesser the number of valence electrons and the farther the valence electrons from the nucleus, the greater is the metallic property. ❖ Metallic property is the ability of the atom to donate electrons. ❖ Metallic property thus increases from top to bottom within the same group and decreases from left to right within the same period. Periodic Property ❖ The Periodic Table was constructed on the basis of trends (periodicity) in chemical properties. ❖ With an understanding of electron configuration, chemists realized that the periodicity of chemical properties could be understood in terms of periodicity in electron configuration. ❖ The Periodic Table works because elements in the same column (group) have the same configuration in their outer shells.  Ionization Energy ❖ The energy required to remove the most loosely held electron from an atom in the gaseous state. example: when lithium loses one electron, it becomes a lithium ion; it still has three protons in its nucleus, but now only two electrons outside the nucleus. ❖ In general, it increases across a row; valence electrons are in the same shell and subject to increasing attraction as the number of protons in the nucleus increases. ❖ it increases going up a column; the valence electrons are in lower principle energy levels, which are closer to the nucleus and feel the nuclear charge more strongly. Electron Affinity ❖ Non-metals, in contrast with metals, tend to gain electrons to form negative ions. The energy in this process is termed as electron affinity. ❖ Electron affinity increases from left to right within the same period and decreases from top to bottom within the same group. Electronegativity ❖ A measure an atom’s attraction for the electrons it shares in a chemical bond with another atom. Ionic Size ❖ When an atom loses or gains an electron it becomes a positively or negatively charged particle ion. ❖ Within the group, the ionic size increases from top to bottom, the ionic size decreases across a period for ions possessing the same electronic configuration  For example, Lesson 4: FORMULA WRITING: Nomenclature of inorganic AuCl Aurous chloride compounds Lewis structure, Resonance structure, Formal charge SnBr2 Stannous bromide PbI2 Plumbous iodide CoSO4 Cobaltous sulfate FORMULA WRITING and NAMING 3. In naming binary compounds containing two non-metals, rule 1 is applied, prefixes are used to denote the number of ❖ A chemical formula represents a compound in terms of chemical elements. symbols to indicate the elements and numbers of atoms present For example, in the compound. PBr5 Phosphorus pentabromide ❖ To write formulas, valances and oxidation numbers are generally CO2 Carbon dioxide used. The valence denotes the number of electrons in the CCl4 Carbon tetrachloride outermost energy level; it also describes the combining power of N2O5 Dinitrogen pentoxide an atom in a compound. The oxidation number refers also to the P4O10 Tetraphosphorus decoxide combining capacity of an atom but it specifies its charge (whether positive or negative). 4. Binary acids are solutions of compounds consisting of Rules in Writing Formulas hydrogen and a nonmetal. They are designated by the prefix 1. The positive ion is written first followed by the negative ion. “hydro” and the suffix “ic” and the word acid is added. 2. The sum of the positive and the negative valences should add For example, up to zero. HCl Hydrochloric acid Examples: HBr Hydrobromic acid Potassium iodide K+ l- KI H2Se Hydroselenic acid Calcium sulfite Ca+2 SO3-2 CaSO3 Aluminum phosphate Al+3 PO4-3 AlPO4 5. Salts of binary acids are formed when the hydrogen of the acid is replaced by other cations. The name of the cation is given first followed by the anion. Naming of Compounds For example, KCl Potassium chloride ❖ The specific name of each compound makes it different from CaS Calcium sulfide another since the identification of the compound by giving the AlBr3 Aluminum bromide name is based on scientific methods. ZnI2 Zinc iodide Naming Ternary Compounds ❖ Ternary compounds consist of three elements- an electropositive element (either hydrogen or a metal) and a polyatomic negative ion. ❖ Ternary oxyacids are compounds made up of hydrogen, Naming Binary Compounds oxygen, and other elements. Suffix terminology is used to differentiate the non-metals between two ternary oxyacids. The 1. Binary compounds contain only two elements in chemical name of the acid containing the non-metal with the lower combination oxidation number ends in “ous: The name of the acid ❖ In naming binary compounds composed of a metal and a containing a non-metal with the higher oxidation number nonmetal, the name of the electropositive element (metal) is carries the suffix “ic”. written first followed by the name of the electronegative element For example, (non-metal) ending in –ide. HNO2 Nitrous acid For example: HNO3 Nitric acid BaO Barium oxide H3PO3 Phosphorous acid MgCl2 Magnesium chloride H3PO4 Phosphoric acid ZnS Zinc sulfide LiH Lithium hydride 2. In naming binary compounds with variable oxidation numbers, there are two systems commonly used. A. Stock system or Roman Numeral System. ❖ The metallic element is written first followed by its oxidation number in Roman numeral enclosed in parenthesis. For example: HgCl2 Mercury (II) chloride FeBr2 Iron (II) bromide CuSO4 Copper (II) sulfate B. Classical or Conventional System. ❖ The suffix “ous” is added to the Latin name of the electropositive metal with the lower oxidation state, and the suffix “ic” applied with the higher oxidation state.  Forming of covalent bond Lesson 5: Intramolecular and Intermolecular Forces of Attraction ❖ A covalent bond is formed by sharing one or more pairs of electrons. OVERVIEW The pair of electrons is shared by both atoms and, at the same time, ❖ Chemical bonding is one of the basic parts of chemistry that fills the valence shell of each atom. explains other concepts such as reactions and molecules. Example: in forming H2, each hydrogen contributes one electron to Without it, chemists wouldn’t be able to give details on why the single bond. atoms are attracted to each other or how products are produced after a chemical reaction has taken place. To know the concept of bonding, one must first know the basics behind the atomic structure. ❖ Chemical bonding occurs when a cation and anion are formed by using a system developed by the International Union of Pure and Applied Chemistry. Many ions have common names that Polarity of covalent bond were in use long before chemists undertook an effort to systematize their naming ❖ All covalent bonds involve sharing of electron pairs; they differ in the ❖ Intermolecular Forces of Attraction are forces of attraction degree of sharing: A nonpolar covalent bond: electrons are shared between unlike charges, partially positive and negative dipoles equally while polar covalent bond shared electrons unequally. The that occur between two molecules. They differ from bonds in degree of sharing depends on the relative electronegativities of the that they are not as strong and occur between two adjacent or bonded atoms. neighboring molecules and not within the molecule itself. They do form, just like bonds, due to the attraction of opposite charges. Metallic bond Forming chemical bonds 1. Chemical bonding is any of the connections or interactions ❖ The force that binds together the atoms of metals is called the that explicate the association of atoms into molecules, ions, metallic bond. crystals, and other stable species that make up the familiar ❖ The atoms of metals are all alike therefore they cannot form ionic substances of the everyday world. Atoms in compounds are held bonds. together by powerful forces of attraction called chemical bonds. ❖ The atom of metallic elements contains only 1 to 3 valence electrons, There are two main types of chemical bonds: ionic bond and therefore these atoms cannot form covalent bonds, with noble gas covalent bond. configurations as they will remain incomplete. ❖ Metallic bonds result from the sharing of a variable number of electrons by a variable number of atoms. INTRAMOLECULAR AND INTERMOLECULAR FORCE ❖ Also the electrostatic attraction between the positively charged atomic nuclei of metal atoms and the delocalised electrons in the ❖ Intramolecular force is any force that binds together the atoms to metal. form a molecule or compound, while intermolecular forces are ❖ Metallic bonding is the principal force holding together the atoms of the forces present between molecules. a metal. Intra – meaning inside ❖ The conduction of heat and electricity that characterize metals is due Inter – between or among to their distinctive kind of bonding. The solid metals behave as if ❖ Chemical bonds are considered to be intramolecular forces. they are composed of positive ions in a sea of electrons (also called These forces are often stronger than intermolecular forces, which mobile electron clouds). are present between atoms or molecules that are not bonded. ❖ In metallic crystal, positively charged ions occupy position in the crystal structure and the loosely bound electrons move freely According to the Lewis model, an atom may lose or gain enough electrons to throughout the crystalline structure attain a filled valence shell and become an ion. An ionic bond is the result of the force of attraction between a cation and an anion. An atom may share electrons with one or more other atoms to acquire a filled valence shell. A covalent bond is the result of the force of attraction between two atoms that share one or more pairs of electrons. Ionic Compounds ❖ According to the Lewis model, an ionic bond is formed by the Electron Sea model for metallic bonding transfer of one or more valence electrons from an atom of lower 1. Lorentz(1923) proposed a model known as the electron gas model or electronegativity to an atom of higher electronegativity. The electron sea model. more electronegative atom gains one or more valence electrons 2. This model is based on the following characteristic properties of and becomes an anion and the less electronegative atom loses metals: electrons and becomes a cation. The compound formed by the Low ionization energies : combination of an anion and a cation is called an ionic ❖ Metals generally have low ionization energies. (low electronegativity) compound. ❖ This implies that the valence electrons of metal atoms are not strongly held by the nucleus. Forming an Ionic Bond ❖ Valence electrons can move freely out of the influence of their kernels ❖ In forming sodium chloride, NaCl, one electron is transferred (atomic orbit/structure minus valence electrons). Thus, metals have from a sodium atom to a chlorine atom: a single-headed curved free mobile electrons. arrow was shown to transfer of one electron:  Metallic bonding Intermolecular Forces of Attraction ❖ Intermolecular forces of attractions are typically noted with a dotted line while bonds are solid lines. Types of intramolecular forces of attraction 1. London Dispersion Forces - Temporary or weak dipole interactions (such as those between nonmetals) in non-polar compounds are called London Dispersion Forces. These are the weakest of the intermolecular forces and exist between all types of molecules, whether ionic or covalent—polar or nonpolar. The more electrons a molecule has, the stronger the London dispersion forces. Concept of bond polarity ❖ Concept of bond polarity is useful in describing the sharing of electrons between atoms. ❖ The greater the difference in electronegativity between two atoms, the more polar the bond. 2. Dipole-Dipole Interactions - Permanent or strong dipole interactions (such as those between nonmetals) in polar compounds are called dipole-dipole interactions. This is definitely the molecular force driving the molecules to be soluble with one another and stay in solution. These forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule. The prerequisite for this type of attraction to exist is partially charged ions— Relative strength of the intramolecular forces 3. Hydrogen Bonds - Hydrogen bonds are a special type of dipole-dipole interaction between the hydrogen atom in a polar bond, such as N-H, O-H, or F-H, and an electronegative O, N, or F atom. The strength of a hydrogen bond is determined by the coulombic interaction between the lone-pair electrons of the electronegative atom and the hydrogen nucleus. Hydrogen bonds are simply very strong and very specific dipoles, which makes them the second strongest intermolecular force attraction, second to ion-dipole interactions.  Lesson 7: Oxidation and reduction Balancing REDOX equation by the Oxidation-Reduction Method OXIDATION NUMBERS The oxidation number method for balancing redox equations involves a series of steps. The equation that follows will illustrate the method. Oxidation numbers or valence numbers are arbitrary numbers based upon rules such as the following: 1. Determine the oxidation number of the atoms in the equation in order to identify atoms undergoing oxidation or reduction. 1. The oxidation number of uncombined element is zero. 2. Balance the total increase and total decrease of oxidation number by placing 2. The common oxidation number of hydrogen in compounds is +1, - 1 for the necessary coefficients, using the smallest numbers possible. hydrides. For oxygen it is -2, but -1 for peroxide. 3. Complete the balancing process by inspection. 3. The common oxidation number for group 7A elements in binary compounds is -1. It varies in tertiary compounds. 4. The common oxidation number for group 1A ions is +1; for group 2A it is LEORA (Loss electron, It OXIDIZE, Reducing Agent) +2, and for group 3A it is +3. 5. The oxidation number of an ion is calculated if the oxidation numbers of all GEROA (Gain Electron, it REDUCE, Oxidizing Agent) the others in the compound are known, since the sum of all oxidation numbers in compound is zero. OXIDATION Oxidation generally involves the combination of a substance with oxygen. For example rusting of iron in which the reaction is 4 Fe + 3 O2 → 2 Fe2O3 Also, when glucose (sugar) in our cells reacts with oxygen, carbon dioxide and water are formed. The reaction is C6H12O6 + O2 → 2 CO2 + 6 H2O OXIDATION - REDUCTION Oxidation- Reduction (REDOX) reactions occur simultaneously. Oxidation involves the loss of electrons which results in an increase in the oxidation number of the atom or ion. Reduction – involves the gain of electrons which results in a decrease in the oxidation number of the atom or ion. Five important types of redox reactionsFive important types of redox reactions 1. Combustion: burning in air. The products of complete combustion of carbon compounds are CO2 and H2O. 2. Respiration: the process by which living organisms use O2 to oxidize carbon-containing compounds to produce CO2 and H2O. The importance of these reaction is not the CO2 produced, but the energy released. 3. Rusting: the oxidation of iron to a mixture of iron oxides 4 Fe(s) + 3 O2 ( g) 2 Fe2 O3 ( s) 4. Bleaching: the oxidation of colored compounds to products which are colorless. 5. Batteries: in all cases, the reaction taking place in a battery is a redox-reaction. In oxidation-reduction reactions, the substance that is oxidized is called the reducing agent, while the substance that is reduced is called the oxidizing agent.  Lesson 8: Chemical reactions Types of Chemical Reactions Typical Decomposition Reactions: Chemical Reactions and a. Hydrates, when heated decompose to yield water and the anhydrous Chemical Equations salt. A hydrate is a salt that contains none or more molecule of water for each formula unit of salt, built right into the crystal structure. ❖ A chemical reaction transforms one or more substances into a set of different BaCl2.2H2O (s) BaCl2 (s) + 2 H2O (g) substances. The substances that enter into a chemical reaction are called reactants and the substances formed are called products. ❖ Chemical equations are representations of chemical reactions in terms of symbols of elements and formulas of compounds involved in the reaction. Symbols commonly used in chemical equations 3. Single Replacement reaction – is one in which a metal replaces another metal ion from a solution or a non-metal replaces a less active nonmetal in a compound. This reaction is also called displacement or substitution. General form: AX + B → BX + A (where A and B are metals) or AX + Y → AY + X (where X and Y are nonmetals) Types of Chemical Reactions 1. Composition reaction – is one in which two or more substances (either elements or compounds) react to form one compound. This reaction is also known as combination, direct union or synthesis. General form: A + B → AB Examples: 1. 2 Na + Cl2 → 2 NaCl 2. CaO + H2O → Ca(OH)2 2. Decomposition reaction – is one in which one compound decomposes to form two or more new substances. Usually heat is necessary to cause this reaction to take place. This is also known as analysis. General form: AB → A + B Examples: 1. 2 KClO3 → 2 KCl + 3 O2 2. 2 HgO → 2 Hg + O2 Examples: 1. Fe + CuSO4 → FeSO4 + Cu 2. Cu + FeSO4 → no reaction 3. Zn + HCl → ZnCl2 + H2  4. Double decomposition reaction – is one wherein two compounds react to form two new compounds.This involves exchange of ion pairs. This reaction is also called exchange reaction or metathesis. General form: AX + BY → AY + BX Examples: 1. 2 HNO3 + Ca(OH)2 → Ca(NO3)2 + 2 H2O 2. Na2SO4 + BaCl2 → BaSO4 + 2 NaCl Balancing an equation ❖ An equation has to be consistent with the fundamental law of nature – The Law of Conservation of Mass. That is, atoms are neither lost nor gained during chemical reactions.