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BIOCHEM REVIEWER CHAPTER 1: MATTER AND ITS PROPERTIES 2. INORGANIC CHEMISTRY – the study of compounds not covered by Introduction to Chemistry organic chemistry; the study of inorganic...

BIOCHEM REVIEWER CHAPTER 1: MATTER AND ITS PROPERTIES 2. INORGANIC CHEMISTRY – the study of compounds not covered by Introduction to Chemistry organic chemistry; the study of inorganic compounds or compounds that do not CHEMISTRY- the study of specific kinds contain a C-H bond. Many inorganic of matter and its changes in compounds are those which contain composition. metals. 3. ANALYTICAL CHEMISTRY – the study of the chemistry of matter and the development of tools used to measure properties of matter. 4. PHYSICAL CHEMISTRY – the branch of chemistry that applies physics to the study of chemistry. Commonly this includes the applications of thermodynamics and quantum mechanics to chemistry. Chemistry: Central Science 5. BIOCHEMISTRY – the study of chemical processes that occur inside of living organisms. DEFINITION OF MATTER - the physical material of the universe; it is anything that has Scientific Method mass and occupies space. CLASSIFICATION OF MATTER Branches of Chemistry PHYSICAL STATE 1. ORGANIC CHEMISTRY – the study SOLIDS – relatively rigid and of carbon and its compounds; the study have fixed shapes and volume of the chemistry of life. ○ Fixed shape and volume. them can be made to LIQUIDS – have fixed volumes conduct electricity. but flow to assume the shape of their containers. ❖ Compounds – contains ○ Fixed volume; no fixed two or more elements and shape has chemical and physical properties that are usually GASES – have neither fixed different from those of the shape nor fixed volumes and elements of which it is expand to completely fill their composed. containers. Organic – derived ○ No fixed shape nor volume from or produced by living organisms and have COMPOSITION carbon-hydrogen covalent - Pure substance bonds. - Mixture Inorganic – derived from nonliving PURE SUBSTANCES – any components, and matter that has a fixed chemical generally have ionic composition and characteristic bonds, lack properties. carbon-hydrogen bonds, ❖ Elements – substances and rarely, if ever, contain that cannot be broken any carbon atoms. down into simpler ones by chemical changes MIXTURES – combinations of Metals – good two or more pure substances in conductors of heat and variable proportions in which the electricity, and are individual substances retain their malleable (they can be identity. hammered into sheets) Homogenous Mixtures and ductile (they can be – exhibit a uniform composition drawn into wire). and appear visually the same Nonmetals – are throughout. usually poor conductors of Heterogeneous heat and electricity, and Mixtures – the composition of a are not malleable or material is not completely uniform ductile. Metalloids – are PROPERTIES OF MATTER intermediate in their PROPERTY – any characteristic that properties. In their physical allows us to recognize a particular type properties, they are more of matter and to distinguish it from other like the nonmetals, but types. under certain circumstances, several of  ATOMS – limit of chemical subdivision for matter. MOLECULES – smallest particle of a pure substance that has the properties of that substance and is capable of a stable independent existence and limit Physical Properties – properties of of physical subdivision for a pure matter that can be observed or substance. measured without attempting to change the composition of the matter being observed. Chemical Properties – properties that matter demonstrate when attempts are made to change the matter into new substances. CLASSIFICATION OF MOLECULES DIATOMIC MOLECULES – contain Intrinsic Properties – are not two atoms. dependent upon how much material is POLYATOMIC MOLECULES – contain present. more than two atoms. HOMOATOMIC MOLECULES – Extrinsic Properties – do depend on contain only one kind of atom. how much material is present. HETEROATOMIC MOLECULES – contain two or more kinds of atoms. CHANGES OF MATTER Physical Changes – do not change CHAPTER 2: ATOMIC THEORY the composition of the substance. Example: cutting of paper. Early Views of Atom Chemical Changes – change in 5th century B.C. matter leads to change in composition. Democritus - All matter is composed Example: burning of magnesium metal. of small, finite particles and called it atomos – “indivisible” or “uncuttable.” Atoms as moving particles that differ in shape and size, could join together. Aristotle – Matter consists of various combinations of the four “elements” — fire, earth, air, and water — and could be infinitely divided. DALTON’S ATOMIC THEORY COMPOSITION OF MATTER 19th century – John Dalton laid for an atomic theory that linked the idea of elements with the idea of atoms. His As scientists developed methods for atomic theory was based on four probing the nature of matter, the postulates: supposedly indivisible atom began to 1. Each element is composed of show signs of a more complex structure, extremely small particles called and today we know that the atom is atoms. composed of subatomic particles. 2. All atoms of a given element are Particles with the same charge repel identical, but the atoms of one one another, whereas particles with element are different from the unlike charges attract one another. atoms of all other elements. CATHODE RAYS AND ELECTRONS 3. Atoms of one element cannot be J.J. Thomson – observed that changed into atoms of a different cathode rays are the same regardless of element by chemical reactions; the identity of the cathode material. atoms are neither created nor Experiments showed that cathode rays destroyed in chemical reactions. are deflected by electric or magnetic fields in a way consistent with their 4. Compounds are formed when being a stream of negative electrical atoms of more than one element charge. combine; a given compound always has the same relative Cathode rays are streams of number and kind of atoms. negatively charged particles and these negatively charged particles are called PROBLEMS WITH DALTON’S ATOMIC electrons. THEORY Matter is composed of tiny indivisible atoms. atoms are divisible and are composed of smaller, subatomic particles called electrons, protons, and neutrons. OIL DROP EXPERIMENTS Robert Millikan – succeeded in Atoms of an element are identical in measuring the charge of an electron by mass. performing the oil-drop experiment in all elements have isotopes – 1909. The electron has a charge of atoms with this same proton number but 1.602 x 10-19 C. different numbers of neutrons atoms of an element do not have to have the same mass. DISCOVERY OF SUBATOMIC PARTICLES  ISOTOPES AND NEUTRONS During the early 1900s, scientists identified several substances that appeared to be new elements, isolating them from radioactive ores. For example, a “new element” produced by the radioactive decay of thorium was initially given the name mesothorium. RADIOACTIVITY However, a more detailed analysis Radioactivity – spontaneous emission showed that mesothorium was of radiation. chemically identical to radium (another Radiation – the emission of energy as decay product), despite having a electromagnetic waves or as moving different atomic mass. subatomic particles, especially high-energy particles that cause Frederick Soddy – discovered an ionization element could have types of atoms with different masses that were chemically Henri Becquerel – discovered that a indistinguishable. And he called these compound of uranium spontaneously isotopes—atoms of the same element emits high-energy radiation in 1896. that differ in mass. Marie and Pierre Curie – began The nucleus was known to contain experiments to identify and isolate the almost all of the mass of an atom, with source of radioactivity in the compound the number of protons only providing at Becquerel’s suggestion. half, or less, of that mass. Different proposals were made to explain what Ernest Rutherford – revealed three constituted the remaining mass, types of radiation: alpha α, beta β, and including the existence of neutral gamma γ. Most of the mass of each particles in the nucleus. As you might gold atom and all of its positive charge, expect, detecting uncharged particles is which he called proton, reside in a very very challenging. small, extremely dense region and called it nucleus. James Chadwick – discovered neutrons: uncharged, subatomic particles with a mass approximately the same as that of protons. ATOMIC MODELS DALTON’S ATOMIC MODEL PLUM-PUDDING MODEL OF ATOM MODERN ATOMIC MODEL MODERN ERA’S VIEW OF ATOM J.J. THOMSON – electrons contribute Following the work of Ernest only a very small fraction of an atom’s Rutherford and his colleagues in the mass; they probably are responsible for early twentieth century, the picture of an equally small fraction of the atom’s atoms consisting of tiny dense nuclei size. The atom consists of a uniform surrounded by lighter and even tinier positive sphere of matter in which the electrons continually moving about the mass is evenly distributed and in which nucleus was well established. This the electrons are embedded like raisins picture was called the planetary model, in a pudding or seeds in a watermelon. since it pictured the atom as a miniature “solar system” with the electrons orbiting the nucleus like planets orbiting the sun. BOHR’S ATOMIC MODEL N. BOHR – attempted to resolve the atomic paradox by ignoring classical electromagnetism’s prediction that the orbiting electron in hydrogen would continuously emit light. Furthermore, he adopted Planck’s idea that energies are quantized and made three postulates: 1. Only orbits of certain radii, corresponding to certain specific NUCLEAR MODEL OF ATOM energies, are permitted for the electron in a hydrogen atom. E. RUTHERFORD – most of the 2. An electron in a permitted orbit is volume of an atom is empty space in in an “allowed” energy state. An which electrons move around the electron in an allowed energy nucleus. state does not radiate energy and, therefore, does not spiral into the nucleus.  3. Energy is emitted or absorbed by have properties of both a particle and a the electron only as the electron wave. Their exact trajectories cannot be changes from one allowed determined. energy state to another. This Radiation appears to have either a energy is emitted or absorbed as wave-like or a particle-like (photon) a photon that has energy E = hν. character. N. BOHR – assumed that the electron orbiting the nucleus would not normally emit any radiation (the stationary state hypothesis), but it would emit or absorb a photon if it moved to a different orbit. L. DE BROGLIE – one of the first people to pay attention to the special behavior of the microscopic world. He asked the question: If electromagnetic radiation can have particle-like character, can electrons and other submicroscopic particles exhibit wavelike character? LIMITATIONS TO BOHR’S MODEL OF ATOM: An electron moving about the nucleus 1. Bohr model explains the line of an atom behaves like a wave and spectrum of the hydrogen atom, it therefore has a wavelength. the cannot explain the spectra of other wavelength of the electron, or of any atoms. other particle, depends on its mass, m, 2. Bohr also avoided the problem of why and on its velocity, v: λ = h/mv the negatively charged electron would not just fall into the positively charged If an electron is viewed as a wave nucleus, by simply assuming it would circling around the nucleus, an integer not happen. number of wavelengths must fit into the 3. There is a problem with describing an orbit for this standing wave behavior to electron merely as a small particle be possible. circling the nucleus because an electron exhibits wavelike properties, a fact that any acceptable model of electronic structure must accommodate. WAVE BEHAVIOR OF MATTER Macroscopic objects act as particles. Microscopic objects (such as electrons)  Arrangement of the elements in a table W. HEISENBERG – considered the based on the periodic law limits of how accurately we can measure In a modern periodic table, elements properties of an electron or other with similar chemical properties are microscopic particles. He determined found in vertical columns called groups that there is a fundamental limit to how or families accurately one can measure both a particle’s position and its momentum simultaneously. The more accurately we measure the momentum of a particle, the less accurately we can determine its position at that time, and vice versa. It is fundamentally impossible to determine simultaneously and exactly GROUP OR FAMILY both the momentum and the position of The U.S. system uses a Roman a particle. numeral and a letter (either A or B) at the top of the column QUANTUM MECHANICAL The IUPAC (International Union of DESCRIPTION OF ATOM Pure and Applied Chemistry)(but not In 1926, Erwin Schrödinger, proposed universally used) system uses only a an equation that incorporates both the number from 1 to 18 wavelike and particle-like behaviors of the electron (Schrödinger equation). 𝐇 ෡ 𝛙 = 𝐄𝛙 His work opened a new approach to dealing with subatomic particles, an approach known as quantum mechanics or wave mechanics. CHAPTER 3: PERIODIC TABLE & PERIODIC TRENDS PERIOD PERIODIC TABLE  Horizontal row of elements arranged Electronic configurations are used to according to increasing atomic numbers classify elements of the periodic table in Numbered from top to bottom numerous ways By the location of the distinguishing electron Last and highest-energy electron found in an element ELEMENT CLASSIFICATION Representative element – element in MODERN PERIODIC TABLE which the distinguishing electron is Elements 58–71 and 90–103 are not found in an s or a p subshell placed in their correct periods but are Transition element – element in located below the main table which the distinguishing electron is found in a d subshell GROUP AND PERIOD Inner-transition element – element in IDENTIFICATION which the distinguishing electron is Elements and the periodic table found in an f subshell Each element belongs to a group and period of the periodic table PERIODIC TRENDS Examples of group and period location specific patterns that are present in the for elements periodic table that illustrate different Calcium, Ca, element 20: group aspects of a certain element, including IIA(2), period 4 its size and its electronic properties. Silver, Ag, element 47: group IB(11), period 5 Periodic trends, arising from the Sulfur, S, element 16: group arrangement of the periodic table, VIA(16), period 3 provide chemists with an invaluable tool to quickly predict an element's ELEMENT CLASSIFICATION properties. These trends exist because Electronic configurations are used to of the similar atomic structure of the classify elements of the periodic table in elements within their respective group numerous ways families or periods, and because of the By the location of the distinguishing periodic nature of the elements. electron By status as a noble gas MAJOR PERIODIC TRENDS By status as a representative, INCLUDE: transition, or inner-transition element 1. Electronegativity By status as a metal, nonmetal, or 2. Ionization energy metalloid 3. Electron Affinity 4. Atomic/ Ionic Radius ELEMENT CLASSIFICATION 5. Metallic Character ELECTRONEGATIVITY  measures an atom's tendency to the electron affinity value, the higher an attract and form bonds with electrons. atom's affinity for electrons. This property exists due to the electronic configuration of atoms. Most atoms follow the octet rule (having the valence, or outer, shell comprise of 8 electrons). Because elements on the left side of the periodic table have less than a half-full valence shell, the energy required to gain electrons is significantly higher ATOMIC RADIUS compared with the energy required to one-half the distance between the lose electrons. nuclei of two atoms (just like a radius is half the diameter of a circle). However, this idea is complicated by the fact that not all atoms are normally bound together in the same way. Some are bound by covalent bonds in molecules, some are attracted to each other in ionic crystals, and others are held in metallic IONIZATION ENERGY crystals. energy required to remove an electron from a neutral atom in its gaseous phase. Conceptually, ionization energy is the opposite of electronegativity. The lower this energy is, the more readily the atom becomes a cation. Therefore, the higher this energy is, the more unlikely it is that the atom becomes a cation. IONIC RADIUS radius of an atom's ion. Although neither atoms nor ions have sharp boundaries, they are sometimes treated as if they were hard spheres with radii such that the sum of ionic radii of the cation and anion gives the distance between the ions in a crystal lattice. ELECTRON AFFINITY Ionic radii are typically given in units of ability of an atom to accept an either picometers (pm) or angstroms electron. Unlike electronegativity, (Å). electron affinity is a quantitative measurement of the energy change that occurs when an electron is added to a neutral gas atom. The more negative  If a metal can form cations with different charges, the positive charge is indicated by a Roman numeral in parentheses following the name of the metal. Ex. Fe2+ iron (II) ion Fe3+ iron (III) ion Cu+ copper (I) ion Cu2+ copper (II) ion METALLIC CHARACTER how readily an atom can lose an An older method still widely used for electron. distinguishing between differently charged ions of a metal uses the endings -ous and -ic added to the root of the element’s Latin name. Ex. Fe2+ ferrous ion Fe3+ ferric ion Cu+ cuprous ion Cu2+ cupric ion CHAPTER 4: CHEMICAL FORMULAS AND NAMING Cations formed from non-metal atoms have names that end in –ium. DEFINITION Ex. NH4 + ammonium ion H3O+ system used in naming substances. hydronium ion are based on the division of substances into categories: organic and ANIONS inorganic compounds. The names of monatomic anions are ORGANIC COMPOUNDS – formed by replacing the ending of the contain carbon and hydrogen, often in name of the element with –ide. combination with oxygen, nitrogen, or Ex. H- hydride ion other elements. O2- oxide ion INORGANIC COMPOUNDS – N3- nitride ion compounds that do not contain carbon, OH- hydroxide ion and not consisting of or deriving from CN- cyanide ion living matter. P3- phosphide ion NAMING IONIC COMPOUNDS Polyatomic anions containing oxygen CATIONS have names ending in either -ate or -ite formed from metal atoms have the and are called oxyanions. The -ate is same name as the metal. used for the most common or Ex. Na+ sodium ion representative oxyanion of an element, Zn2+ zinc ion and -ite is used for an oxyanion that has Al3+ aluminum ion the same charge but one O atom fewer. Prefixes are used when the series of oxyanions of an element extends to four K+ + OH- → KOH potassium hydroxide members, as with the halogens. The Ca2+ + OH- → Ca(OH)2 calcium prefix per- indicates one more O atom hydroxide than the oxyanion ending in -ate; hypo- Al3+ + OH- → Al(OH)3 aluminum indicates one O atom less than the hydroxide oxyanion ending in –ite. Examples NO3 - nitrate ion ACIDS NO2 - nitrite ion Acids, however, are named slightly SO4 2- sulfate ion differently based upon whether or not SO3 2- sulfite ion they are oxyacids or non-oxyacids. ClO4 - perchlorate ion Oxyacids are acids that have oxygen in ClO3 - chlorate ion the formula. Non-oxyacids are acids that ClO2 - chlorite ion do not contain oxygen in the formula. ClO- hypochlorite ion Non-oxyacids – acids containing anions whose names end in Anions derived by adding H+ to an -ide are named by changing the -ide oxyanion are named by adding as a ending to -ic, adding the prefix hydro- to prefix the word hydrogen or bi- or this anion name, and then following with dihydrogen, as appropriate. the word acid. Ex. CO3 2- carbonate ion Examples: H+ + Cl- → HCl HCO3 - hydrogen carbonate ion hydrochloric acid PO4 3- phosphate ion H+ + S2- → H2S hydrosulfuric acid HPO4 2- hydrogen phosphate ion H2PO4 - dihydrogen phosphate ion Oxyacids – Acids containing anions whose names end in -ate or -ite IONIC COMPOUNDS are named by changing -ate to -ic and Names of ionic compounds consist of -ite to -ous and then adding the word the cation name followed by the anion acid. Prefixes in the anion name are name. retained in the name of the acid. Ex. Ca2+ + Cl- → CaCl2 calcium Examples: H+ + ClO4 - → HClO4 chloride perchloric acid Al3+ + NO3 - → Al(NO3)3 aluminum H+ + ClO3 - → HClO3 chloric acid nitrate H+ + ClO2 - → HClO2 chlorous acid Cu2+ + ClO4 - → Cu(ClO4 )2 copper (II) H+ + ClO- → HClO hypochlorous acid perchlorate or cupric perchlorate NAMING BINARY COMPOUNDS NAMING ACIDS AND BASES The procedures used for naming BASES binary (two-element) molecular Naming bases does not necessitate compounds are similar to those used for any change in the rules. Bases always naming ionic compounds. have OH- as the anion with some metal 1. The name of the element farther to as the cation. the left in the periodic table (closest to Ex. Na+ + OH- → NaOH sodium the metals) is usually written first. An hydroxide exception occurs when the compound contains oxygen and chlorine, bromine, A polar covalent bond results when the or iodine (any halogen except fluorine), atoms differ in electronegativity. in which case oxygen is written last. 2. If both elements are in the same An ionic bond results when the group, the one closer to the bottom of electronegativity difference is very large, the table is named first. meaning that the electron density is 3. The name of the second element is shifted far toward the element with given an -ide ending. highest electronegativity. 4. Greek prefixes indicate the number of atoms of each element. (Exception: The The greater the difference in prefix mono- is never used with the first electronegativity between two atoms, element.) When the prefix ends in a or o the more polar their bond. and the name of the second element begins with a vowel, the a or o of the DIPOLE – two electrical charges of prefix is often dropped. equal magnitude but opposite signs are separated by a distance. Examples: Cl2O dichlorine monoxide N2O4 dinitrogen tetroxide DIPOLE MOMENT – quantitative NF3 nitrogen trifluoride measure of the magnitude of a dipole. P4S10 tetraphosphorus decasulfide CHAPTER 5: BOND POLARITY AND MOLECULAR GEOMETRY BOND POLARITY a measure of how equally or unequally the electrons in any covalent bond are shared. NONPOLAR COVALENT BOND – one in which the electrons are shared equally, as in Cl2 and N2. POLAR COVALENT BOND – one of the atoms exerts a greater attraction for the bonding electrons than the other. ELECTRONEGATIVITY AND BOND POLARITY A nonpolar covalent bond results when the electronegativities of the bonded atoms are equal.  that minimizes the repulsions among them. The shapes of different ABn molecules or ions depend on the number of electron domains surrounding the central atom. ELECTRON DOMAIN-GEOMETRY – arrangement of electron domains about the central atom of an ABn molecule or ion. VSPER THEORY AND MOLECULAR GEOMETRY MOLECULAR GEOMETRY – arrangement of only the atoms in a MOLECULAR SHAPES molecule or ion—any nonbonding pairs Lewis structures, however, do not in the molecule are not part of the indicate the shapes of molecules; they description of the molecular geometry. simply show the number and types of bonds and are drawn with the atoms all in the same plane. BOND ANGLES – angles made by the lines joining the nuclei of the atoms in the molecule. BOND LENGTH – distance between the center of two bonded atoms. The bond angles of a molecule, together with the bond lengths, define the shape and size of the molecule. VSEPR THEORY – stands for Valence Shell Electron Pair Repulsion Theory, based on the idea that electron domains are negatively charged and therefore repel one another. ELECTRON DOMAIN – a region about a central atom in which an electron pair is concentrated. Each nonbonding pair, single bond, or multiple bond produces a single electron domain around the central atom in a molecule. The best arrangement of a given number of electron domains is the one  VSEPR steps in predicting the shapes of molecules or ions: 1. Draw the Lewis structure of the molecule or ion, and count the number of electron domains around the central atom. Each nonbonding electron pair, each single bond, each double bond, and each triple bond counts as one electron domain. 2. Determine the electron-domain geometry by arranging the electron domains about the central atom so that the repulsions among them are minimized. 3. Use the arrangement of the bonded atoms to determine the molecular INTERMOLECULAR FORCES geometry. the forces which mediate interaction between molecules, including forces of attraction or repulsion which act between molecules and other types of neighboring particles, e.g., atoms or ions. Intermolecular forces are weak relative to intramolecular forces – the forces which hold a molecule together. INTERMOLECULAR FORCES VS. INTRAMOLECULAR FORCES Intramolecular forces – the forces that hold atoms together within a molecule. Intermolecular forces – forces that exist between molecules. TYPES INTRAMOLECULAR FORCES Metallic Bond – the strong Ionic Bond – formed by the complete electrostatic force of attraction between transfer of valence electron(s) between metal cations/atoms and delocalized atoms. It is a type of chemical bond that electrons in the metallic lattice of a generates two oppositely charged ions. metallic substance (e.g. the elements in - Give and take group 1 and 2 of the periodic table). This type of bonding only exists in metallic substances (as they consist of metal cations arranged in a regular lattice structure). as charge density increases the Covalent Bond – formed between strength of the metallic substance atoms that have similar increases electronegativities—the affinity or desire for electrons. Because both atoms have similar affinity for electrons and neither has a tendency to donate them, they share electrons in order to achieve octet configuration and become more stable. - Sharing Non Polar Covalent Bond – formed between same atoms or atoms RELATIVE STRENGTH OF with very similar electronegativities—the INTRAMOLECULAR FORCES difference in electronegativity between bonded atoms is less than 0.5. - Mostly Diatomic Atoms TYPES INTERMOLECULAR FORCES Polar Covalent Bond – formed when atoms of slightly different London Dispersion Forces – the electronegativities share electrons. The weakest of the intermolecular forces and difference in electronegativity between a temporary attractive force due to the bonded atoms is between 0.5 and 1.7 formation of temporary dipoles in a nonpolar molecule. When the electrons in two adjacent atoms are displaced in such a way that atoms get some temporary dipoles, they attract each other through the London dispersion force. Ion-Dipole Forces – exist between an ion and a polar molecule. Cations are attracted to the negative end of a dipole, and anions are attracted to the positive end. The magnitude of the attraction increases as either the ionic charge or Dipole-dipole – forces occur when the the magnitude of the dipole moment partially positively charged part of a increases. 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 INTERMOLECULAR FORCES Hydrogen Bond – a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen, nitrogen, or fluorine atom. The partially positive end of hydrogen is attracted to the partially HOW FORCES OF ATTRACTION negative end of the oxygen, nitrogen, or AFFECT PROPERTIES OF fluorine of another molecule. Hydrogen COMPOUNDS? bonding is a relatively strong force of attraction between molecules, and considerable energy is required to break hydrogen bonds. CHAPTER 1.1: France during the French MEASUREMENTS Revolution. The original standards for the meter and the DEFINITION kilogram were adopted there in - provide the macroscopic 1799 and eventually by other information that is the basis of countries. most of the hypotheses, theories, and laws that describe the - Length – measurement or extent behavior of matter and energy in of something from end to end. both the macroscopic and microscopic domains of - Mass – measure of the amount chemistry of material in an object. - provides three kinds of information: - Temperature – measure of the - size or magnitude of the hotness or coldness of an object measurement (a number) is a physical property that - standard of comparison for the determines the direction of heat measurement (a unit) flow. - indication of the uncertainty of the measurement - Time – indefinite continued progress of existence and events MEASUREMENT UNITS in the past, present, and future - Consist of a number and an regarded as a whole identifying unit - Examples - Gallons, Celsius, and DERIVED SI UNITS Fahrenheit - obtained by multiplication or - Based on units agreed on by division of one or more of the those making and using the base units measurements - Volume – measure of the amount of space occupied by an object SI UNITS OR INTERNATIONAL - Density – defined as the amount SYSTEM OF UNITS of mass in a unit volume of a substance - standards for units that are fixed by international agreement METRIC SYSTEM SI BASE UNITS - Initial units of the metric system, which eventually evolved into the SI system, were established in  - Converting Celsius to Kelvin: COMMON METRIC UNITS USED TEMPERATURE UNITS AND CONVERSION ACCURACY AND PRECISION ACCURACY refers to how closely individual measurements agree with the correct, or “true,” value. PRECISION TEMPERATURE UNITS AND a measure of how closely individual CONVERSION measurements agree with one another - Converting Fahrenheit to Celsius: SCIENTIFIC NOTATION - Provides a convenient way to express very large or very small - Converting Celsius to numbers Fahrenheit: - Numbers are represented in the form of M×10n - - Nonexponential term M is a number between 1 to 10 (but not equal to 10) written with a decimal - Exponential term is a 10 - Converting Kelvin to Celsius: raised to a whole number exponent n - n may be positive or negative 10010000000 = 1.001 × 1010 3. If a number is less than 1, then 0.0000001001 = 1.001 × 10−7 only the zeros that are at the end of the number and the zeros that are between nonzero digits are ADDITION AND SUBTRACTION significant. - To add and subtract numbers in scientific notation, make the 0.4800 exponents the same first before This number has 4 significant figures doing the operation with the integers M1 and M2 0.0000910 This number has 3 significant figures (7.4×103) + (2.1×103) = ? 4. If a number is greater than 1, = (7.4+2.1)×103 then all the zeros written to the = 9.5×103 right of the decimal point count as significant figures provided that a Rewrite the scientific notations so the decimal point is present. exponents will be the same. 1530.900 This number has 7 significant figures ((2.43×103) + (3.1×102) = ? (2.43×103) + (0.31×103) = ? 2.060 = (2.43 + 0.31)×103 = 2.74×103 This number has 4 significant figures - all digits of a measured quantity, 5. For numbers that do not contain including the uncertain one decimal points, the trailing zeros (that is, zeros after the last - Determination of significant nonzero digit) may or may not be figures: significant, leading to ambiguity. This can be resolved with 1. In any measurement that is scientific notation, or to place a properly reported, all nonzero decimal point at the end. digits are significant. 500 35612 This is an ambiguous case as there is This number has 5 significant figures no decimal point. This can be rewritten into scientific numbers to eliminate 1398 ambiguity. This number has 4 significant figures OPERATIONS INVOLVING 2. Zeros between nonzero digits are SIGNIFICANT FIGURES always significant. - In addition and subtraction, the 398705 answer must be expressed in This number has 6 significant figures terms of the number with the least decimal places. 101010101 This number has 9 significant figures 89.332 + 1.1 = 90.432 = 90.4  2.097 − 0.12 0.0833 pound (lb.). What is this = 1.977 mass in milligrams (mg)? (1 lb. = = 1.98 453.6 g) OPERATIONS INVOLVING SIGNIFICANT FIGURES In multiplication and division, CHAPTER 2.1 ATOMIC the number of significant figures STRUCTURE in the final product or quotient is determined by the original ATOMS – DEFINITION number that has the smallest - the smallest constituent unit of number of significant figures. ordinary matter that has the properties of a chemical element. 89.332 × 5.1 = 455.5932 = 4.6 x 102 Every solid, liquid, gas, and plasma is composed of neutral or 2.097 ÷ 1.12 =1.87232… = 1.87 ionized atoms. Atoms are extremely small; typical sizes are Dimensional Analysis around 100 picometers (a - The procedure we use to convert ten-billionth of a meter, in the between units in solving short scale). chemistry problems is called dimensional analysis (also called the factor-label method). given unit×conversion factor=desired unit STRUCTURE OF ATOM EXAMPLE PROBLEMS - Convert 12.9 cm to inches. The conversion factor to be used here is 2.54 cm = 1 inch. ATOMIC NUMBER & MASS NUMBER - ATOMIC NUMBER (Z) – number - A person’s average daily intake of of protons in an atom of any glucose (a form of sugar) is particular element. Because an  atom has no net electrical Anions – negatively charged charge, the number of electrons it ions. contains must equal the number O2- Cl- of protons. - Metal atoms tend to lose Z = no. of protons electrons to form cations and non-metal atoms tend to gain - MASS NUMBER (A) – number of electrons to form anions. Thus, protons plus neutrons in the ionic compounds tend to be atom. composed of metals bonded with A = no. of protons + no. of neutrons non-metals. - POLYATOMIC IONS – consist of atoms joined as in a molecule, but carrying a net positive or negative charge. - IONIC COMPOUNDS – a ISOTOPES compound made up of cations - ISOTOPES – atoms with identical and anions. atomic numbers but different Ionic compounds are generally mass numbers (that is, same combinations of metals and number of protons but different non-metals, as in NaCl. In contrast, numbers of neutrons). molecular compounds are generally composed of non-metals only, as in H2O. CHEMICAL FORMULA – a way of expressing information about the proportions of atoms that constitute a particular chemical compound, using a single line of chemical element symbols, numbers, and sometimes also other symbols, such as parentheses, dashes, brackets, commas and plus (+) and minus (−) signs. - MOLECULES – the smallest particle in a chemical element or IONS AND IONIC COMPOUNDS compound that has the chemical - IONS – formed when electrons properties of that element or are removed from or added to an compound and is made up of atom. atoms that are held together by Cations – positively charged chemical bonds. ions. Ca2+ Li+  - MOLECULAR COMPOUNDS – LIMITATIONS TO NIELS BOHR’S composed of molecules contain MODEL OF ATOM more than one type of atom. - Bohr model explains the line spectrum of the hydrogen atom, it - STRUCTURAL FORMULAS – cannot explain the spectra of shows which atoms are attached other atoms. to which and usually does not - Bohr also avoided the problem of depict the actual geometry of the why the negatively charged molecule, that is, the actual electron would not just fall into angles at which atoms are joined. the positively charged nucleus, The atoms are represented by by simply assuming it would not their chemical symbols, and lines happen. are used to represent the bonds - There is a problem with that hold the atoms together. describing an electron merely as a small particle circling the Structural formula representation: nucleus because an electron a.perspective drawing exhibits wave-like properties, a b.ball-and-stick model fact that any acceptable model of c.space-filling model electronic structure must accommodate. QUANTUM MECHANICAL MODEL - In 1926, Erwin Schrödinger, proposed an equation that incorporates both the wave-like and particle-like behaviors of the CHAPTER 3.1: QUANTUM electron. NUMBERS & ELECTRON - His work opened a new approach CONFIGURATION to dealing with subatomic particles, an approach known as Bohr’s Theory of Atomic Structure quantum mechanics or wave - Bohr proposed that the electron mechanics. in a hydrogen atom moved in any - Precise paths of electrons one of a series of circular orbits moving around the nucleus around the nucleus: cannot be determined accurately - Electron could change - Location and energy of electrons orbits only by absorbing or around a nucleus is specified releasing energy using shell, subshell, and orbital - Absorption moves electrons to higher-energy orbits - Release moves electrons to lower-energy orbits - ORBITS – in the Bohr model, which visualizes the electron moving in a physical orbit, like a planet around a star.  occupies in an atom. It can have positive integral values 1, 2, 3 and so on. As n increases, the orbital becomes larger, and the electron spends more time farther from the nucleus. An increase in n also means that the electron has a higher energy and is therefore less tightly bound to the nucleus. - ORBITALS – solution to Schrödinger’s equation for the hydrogen atom yields a set of wave functions, and has a characteristic shape and energy. - ANGULAR MOMENTUM QUANTUM NUMBER (l) – quantum number distinguishing the different shapes of orbitals; QUANTUM NUMBERS - It is also a measure of the orbital angular momentum. It can have - QUANTUM NUMBERS – integral values from 0 to (n – 1) describe values of conserved for each value of n. quantities in the dynamics of a quantum system. In the case of - This quantum number defines the electrons, the quantum numbers shape of the orbital. The value of can be defined as "the sets of l for a particular orbital is numerical values which give generally designated by the acceptable solutions to the letters s (sharp), p (principal), d Schrödinger wave equation for (diffuse), and f (fundamental), the hydrogen atom". corresponding to l values of 0, 1, 2, and 3: - Bohr model introduced a single quantum number, n, to describe an orbit while the quantum mechanical model uses three quantum numbers, n, l, and ml, ANGULAR MOMENTUM QUANTUM which result naturally from the NUMBER (l) mathematics used to describe an orbital. - PRINCIPAL QUANTUM NUMBER (n) – quantum number specifying the shell an electron  RESTRICTIONS ON POSSIBLE VALUES OF n, l AND ml - The shell with principal quantum number n consists of exactly n subshells. - MAGNETIC QUANTUM - Each subshell consists of a NUMBER (ml) – quantum specific number of orbitals. number signifying the orientation - The total number of orbitals in a of an atomic orbital around the shell is n2, where n is the nucleus; principal quantum number of the shell. - orbitals having different values of ml but the same subshell value of l have the same energy (are degenerate), but this degeneracy can be removed by application of an external magnetic field. - It can have integral values between -l and l, including zero. This quantum number describes SPIN QUANTUM NUMBER (ms) – the orientation of the orbital in number specifying the electron spin space. direction. It may have values of +1/2 or -1/2. ELECTRON SHELL – collection of orbitals with the same value of n. ELECTRON SUBSHELL – set of orbitals that have the same n and l values. SUMMARY - PAULI EXCLUSION PRINCIPLE – No two electrons in an atom can have the same set of four quantum numbers n, l, ml, and ms. - HUND’S RULE – states that for degenerate orbitals, the lowest energy is attained when the number of electrons having the ELECTRON CONFIGURATION same spin is maximized. - VALENCE SHELL – outermost - AUFBAU PRINCIPLE – states (highest-energy) shell of an that, hypothetically, electrons element that contains electrons. orbiting one or more atoms fill the Atoms with the same number of lowest available energy levels electrons in the valence shell before filling higher levels (e.g., have similar elemental chemical 1s before 2s). In this way, the properties. electrons of an atom, molecule, or ion harmonize into the most - Detailed arrangement of stable electron configuration electrons in atoms indicated by a possible. specific notation, 1s22s22p4, etc. - When it is remembered that each - Occupied subshells are indicated orbital of a subshell can hold a by their identifying number and maximum of two electrons and letter such as 2s and the number that Hund's rule and the Pauli of electrons in the subshell is exclusion principle are followed, it indicated by the superscript on results in the following filling the letter. order for the first 10 electrons: - Thus, in 1s22s22p4, the 2s2 notation indicates that the 2s subshell contains two electrons. SUBSHELL FILLING ORDER - Electrons will fill subshells in the - Maximum number of electrons order of increasing energy of the each subshell can hold must be subshells remembered - A 1s subshell will fill before a 2s - s subshells can hold 2 subshell - p subshells can hold 6 - Order must obey Aufbau - d subshells can hold 10 principle, Hund's rule and the - f subshells can hold 14 Pauli exclusion principles WRITING ELECTRON GILBERT N. LEWIS – suggested a CONFIGURATIONS simple way of showing the valence Magnesium, Mg, 12 electrons: electrons in an atom and tracking them 1s22s22p63s2 during bond formation. Silicon, Si, 14 electrons: 1s22s22p63s23p2 LEWIS SYMBOL – for an element Iron, Fe, 26 electrons: consists of the element’s chemical 1s22s22p63s23p64s23d6 symbol plus a dot for each valence Gallium, Ga, 31 electrons: electron. 1s22s22p63s23p64s23d104p1 OCTET RULE – atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons. An octet of electrons consists of full s NOBLE GASES ELECTRON and p subshells in an atom. In a Lewis CONFIGURATIONS symbol, an octet is shown as four pairs - With the exception of helium, all of valence electrons arranged around noble gases (Group VIIIA/Column the element symbol. 18) have electronic configurations that end with completely filled s and p subshells of the highest occupied shell - Can be used to write electronic configurations in an abbreviated form in which the noble gas IONIC BONDING symbol enclosed in brackets is IONIC BOND – bond between used to represent all electrons oppositely charged ions. The found in the noble gas ions are formed from atoms by configuration. transfer of one or more electrons. The principal reason ionic compounds are stable is the attraction between ions of CHAPTER 4.1: CHEMICAL opposite charge. This attraction BONDING draws the ions together, releasing energy and causing BASIC CONCEPTS the ions to form a solid array, or lattice. CHEMICAL BOND – strong attractive force that exists between atoms in a In forming ions, transition metals lose molecule. the valence-shell s electrons first, then as many d electrons as required to VALENCE ELECTRONS – are those in reach the charge of the ion. the outermost occupied shell which are involved in chemical bonding. LEWIS DOT STRUCTURE AND OCTET RULE COVALENT BONDING COVALENT BOND – bond formed between two or more atoms by a sharing of electrons. Each shared electron pair as a line and any unshared electron pairs (also called lone pairs of nonbonding pairs) as dots in a Lewis dot structure. SINGLE BOND – shared electron pair constitutes a single covalent bond. DOUBLE BOND – two electron pairs are shared by two atoms, two lines are drawn in the Lewis structure. DRAWING LEWIS STRUCTURE TRIPLE BOND – corresponds to the 1.Sum the valence electrons from all sharing of three pairs of electrons. atoms, taking into account overall charge. 2.Write the symbols for the atoms, show which atoms are attached to which, and connect them with a single bond (a line, representing two electrons). 3.Complete the octets around all the atoms bonded to the central atom. 4.Place any leftover electrons on the LEWIS STRUCTURE central atom. a very simplified representation of the 5.If there are not enough electrons to valence shell electrons in a molecule. It give the central atom an octet, try is used to show how the electrons are multiple bonds. arranged around individual atoms in a molecule. Electrons are shown as "dots" or for bonding electrons as a line between the two atoms. The goal is to obtain the "best" electron configuration, i.e. the octet rule and formal charges need to be satisfied. RESONANCE STRUCTURE placement of the atoms in these two alternative but completely equivalent Lewis structures is the same, but the placement of the electrons is different. EXCEPTIONS TO THE OCTET RULE 1. Molecules and polyatomic ions containing an odd number of electrons 2. Molecules and polyatomic ions in which an atom has fewer than an octet of valence electrons 3. Molecules and polyatomic ions in which an atom has more than an octet of valence electrons  - Symbols indicating the physical state of each reactant and product: (g) = gas; (l) = liquid; (s) = solid; (aq) = aqueous (dissolved in aqueous [water] solution); ↓ = precipitate (during precipitation reaction) - Sometimes symbols that represent the conditions under which the reaction proceeds appear above or below the CHAPTER 5.1: CHEMICAL reaction arrow: REACTIONS Δ = heat; hν = light A (s) + B (l) → C (g) + D (aq) CHEMICAL REACTIONS AND BALANCING CHEMICAL EQUATIONS EQUATIONS - Write down your given equation. - CHEMICAL REACTIONS – Example: processes in which one or more C3H8 + O2 → H2O + CO2 substances are converted into - Write down the number of atoms other substances; also called per each element that you have chemical changes. on each side of the equation. Look at the subscripts next to - CHEMICAL EQUATIONS – each atom to find the number of representation of chemical atoms in the equation. reactions using the chemical Left side: 3 carbon, 8 hydrogen and 2 formulas of the reactants and oxygen products; a balanced chemical Right side: 1 carbon, 2 hydrogen and 3 equation contains equal numbers oxygen of atoms of each element on both - Always leave hydrogen and sides of the equation. oxygen for last. - If you have more than one Example: element left to balance: select the 2 H2 + O2 → 2 H2O element that appears in only a - Reads as + sign as “reacts with” single molecule of reactants and and the arrow as “produces”. The in only a single molecule of chemical formulas to the left of products. This means that you the arrow represent the starting will need to balance the carbon substances, called reactants. The atoms first. chemical formulas to the right of - Add a coefficient to the single the arrow represent substances carbon atom on the right of the produced in the reaction, called equation to balance it with the 3 products. The numbers in front of carbon atoms on the left of the the formulas, called coefficients, equation. Example: indicate the relative numbers of C3H8 + O2 → H2O + 3 CO2 molecules of each kind involved - Balance the hydrogen atoms in the reaction. next. You have 8 on the left side. - Indicating the states of reactants So you'll need 8 on the right side. and products C3H8 + O2 → 4 H2O + 3 CO2 - Balance the oxygen atoms. C3H8 + 5 O2 → 4 H2O + 3 CO2  metal hydroxides → metal oxides TYPES OF CHEMICAL REACTIONS + H2O Most chemical reactions can be Cu(OH)2 (s) → CuO (s) + H2O (g) classified as being: 1. Combination (synthesis) binary compounds → 2 elements 2. Decomposition (analysis) 2 NaN3 (s) → 2 Na (s) + 3 N2 (g) 3. Combustion 4. Single Replacement (substitution) oxyacids → non-metal oxide + H2O 5. Double Replacement 2 HNO3 (aq) → H2O (l) + N2O5 (g) (metathesis) 6. Redox (reduction-oxidation) COMBUSTION – organic hydrocarbons, CxHy, burn in oxygen to form carbon - COMBINATION – two or more dioxide and water vapor (which also substances combine to form a produce heat), incomplete combustion single, more complex compound, may result in carbon monoxide and many elements react with one carbon soot and has the general another in this fashion to form equation of: compounds and has the general CxHy + O2 (g) → __CO2 (g) + __H2O equation of: (g) A + X → AX C3H8 + 5 O2 (g) → 4 H2O (g) + 3 - A combination reaction between CO2 (g) a metal and non-metal products - Most commonly, C-H bond siya produces an ionic solid. Example: and ang result niya ay mostly 2 Mg (s) + O2 (g) → 2 MgO (s) CO2 and H2O - If 2 or more simple compounds react, they follow common - SINGLE REPLACEMENT – a patterns. Ex. metal replaces another metal in a non-metal oxides + water → compound or a non-metal acids replaces another non-metal in a CO2 (g) + H2O (l) → H2CO3 (aq) compound and has the general metal oxides + water → bases equation of: CaO (s) + H2O (l) → Ca(OH)2 (aq) AB + C → CB + A or AB + D → AD + B - DECOMPOSITION – one - Replacement of a metal by a substance breaks down into more active metal example: simpler compounds or elements, - molecular equation: often occurs when compounds CuSO4 (aq) + Mg (s) → MgSO4 (aq) + are heated or electricity is added Cu (s) and has the general equation of: - ionic equation: AX → A + X Cu2+ + SO42- + Mg → Mg2+ + SO42- - Common reactivity patterns of + Cu decomposition reactions - omission of spectator examples: ions: Cu2+ + Mg → Mg2+ + Cu metal chlorates → metal - Replacement of a non-metal by a chlorides + O2 more active nonmetal example: 2 KClO3 → 2 KCl + 3 O2 2 NaI (aq) + Cl2 (g) → 2 NaCl (aq) + I2 (s) metal carbonates → metal oxides + CO2 - DOUBLE DISPLACEMENT – the Na2CO3 → Na2O + CO2 metals in two aqueous  compounds switch places and has the general equation of: - Reduction: AB + CD → CB + AD To lose oxygen AgNO3 (aq) + NaCl (aq) → NaNO3 To combine with hydrogen (aq) + AgCl (s) To gain electrons To decrease in oxidation number REVERSIBLE AND IRREVERSIBLE REACTIONS - IRREVERSIBLE REACTION – reaction in which the products are not available to react to form the initial reactants because: a.precipitate forms b.gas forms c.liquid forms d.product is removed - REVERSIBLE REACTION – reaction in which the products remain available to react to form the initial reactants. REVERSIBLE AND IRREVERSIBLE REACTIONS IRREVERSIBLE REACTION REVERSIBLE REACTION - REDOX – a combination of two words, reduction and oxidation. These two words have multiple meanings when applied to chemical reactions - Oxidation: To combine with oxygen To lose hydrogen To lose electrons To increase in oxidation number