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Final Exam Study Guide Che 106

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Final Exam Study Guide Che 106 CHE 106 - M001

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Chapters 1-8 for Doyle's class.
General Chemistry 1
Dr. Doyle
Study Guide
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This 79 page Study Guide was uploaded by Bianca on Wednesday December 23, 2015. The Study Guide belongs to CHE 106 - M001 at Syracuse University taught by Dr. Doyle in Summer 2015. Since its upload, it has received 37 views. For similar materials see General Chemistry 1 in Chemistry at Syracuse University.

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Date Created: 12/23/15
Introduction: Matter and Measurement 12/29/15 8:28 PM 1.1 The Study of Chemistry Chemistry is the study of the properties and behavior of matter. Matter is the physical material of the universe; it is anything that has mass and occupies space. A property is any characteristic that allows us to recognize a particular type of matter and to distinguish it from other types. Elements are matter comprised of combinations of only about 100 substances. Atoms are the almost infinitesimally small building blocks of matter. An element is composed of a unique kind of atom. The properties of matter relate to both the kinds of atoms the matter contains (composition) and the arrangements of these atoms (structure). Molecules are two or more atoms joined in specific shape. 1.2 Classifications of Matter Matter is characterized by its physical state (liquid, solid, gas) and composition (element, compound or mixture). States of matter- liquid, solid, gas Gas (vapor)- has no fixed volume or shape and uniformly fills it container. Liquid- has distinct volume independent of its container. Solid- definite shape and definite volume. A gas can be compressed to occupy a smaller volume but a liquid and solid cannot. Changes in temperature and/or pressure can lead to conversion from one state of matter to another. A pure substance is matter that has distinct properties and a composition that does not vary from sample to sample. All substances are either elements or compounds. Elements are substances that cannot be decomposed into simpler substances. Compounds are substances composed of two or more elements- contains two or more kinds of atoms. Mixtures are combinations of two or more substances in which each substance retains its chemical identity. Elements oxygen (O), carbon (C), and hydrogen (H) account for over 90% of the mass of the human body. In the periodic table- elements are arranged in columns so closely related elements are grouped together. Elements interact with other elements to form compounds. Law of constant composition/law of definite proportions- elemental composition of a compound is always the same. A pure compound has the same composition and properties under the same conditions regardless of its source (Joseph Louis Proust). Each substance in a mixture retains its chemical identity and properties. Substances that make up a mixture are called the components of the mixture. Heterogeneous mixtures are mixtures that do not have the same composition, properties and appearance throughout. Homogenous mixtures are mixtures that are uniform throughout. Homogenous mixtures are also called solutions. Solutions can be solids, liquids or gases. Classification of matter: All pure matter is classified ultimately as either an element or a compound. 1.3 Properties of Matter Physical properties can be observed with changing the identity and composition of the substance. Ex- color, odor, density, melting point, boiling point, and hardness. Chemical properties describe the way a substance may change, or react, to form other substances. Ex- flammability. Intensive properties do not depend on the amount of sample being examined and are particularly useful in chemistry because many intensive properties can be useful to identity substances. Ex- temperature and melting point. Extensive properties depend on the amount of the substance present. Ex- mass and volume. Physical change- substance changes its physical appearance but not its composition. Ex- evaporation of water, changes from liquid to gas state. Changes of state are physical changes. Chemical change/ reaction- a substance is transformed into a chemically different substance. Ex- when hydrogen burns in air. A method used to separate components of a homogeneous mixture is distillation, a process that depends on the different abilities of substances to form gases. Chromatography- a technique used to separate mixtures by adhering to the substances of solids. Filtration- a mixture of a solid and a liquid is poured through filter paper. The liquid passes through the paper while the solid remains on the paper. Quantitative properties of matter are those associated with numbers. The units used for scientific measurements are those of the metric system. The scientific method: A hypothesis is a tentative explanation to be tested. A theory is a well-substantiated explanation of some aspect of the natural world, based on a body of facts that have been repeatedly confirmed through observation and experiment. Scientific law- natural behavior repeatedly occurring under all different sorts of conditions. 1.4 Units of Measurement SI units contains 7 base units from which all other units are derived. Ex- the basic units for length (meter), mass (kilogram) and temperature (kelvin). Mass (kg) is a measure of the amount of material in an object. Temperature is a physical property that measures the hotness or coldness of an object by its direction of heat flow. Temperature is measured primarily in Kelvin and Celsius. Celsius has a freezing point of 0°C in water and a boiling point of 100°C at sea level. The Kevin (K) scale is the SI unit used. Zero on the Kelvin scale is the lowest attainable temperature- referred to as absolute zero. According to the Fahrenheit scale water freezes at 32°F and boils at 212°F. A derived unit is obtained by multiplication or division of one or more of the base units. Ex- speed =m/s Volume of a cube is its length cubed (cubic meter). Density (g/cm^3) or (g/mL) is defined as the amount of mass in a unit of volume of a substance: Density= mass/volume Numbers obtained by measurement is always inexact, uncertainty always exists in measured quantities. 1.5 Uncertainty in Measurement Precision- how closely individual measurements agree with one another. Accuracy- how closely individual measurements agree with the correct value. Rules for significant figures: All nonzero numbers are significant Zeros between nonzero numbers are significant Zeros at the beginning of a number are never significant- they only indicate position of decimal point. Zeros at the end of a number are significant if the number contains a decimal. Ex- .0200g has 3 significant figures. If zeros at the end of a number are significant the number will be displayed in exponential notation. Ex- 1.030x10^4g has 4 sig figs. Rules for significant figures in calculations: For addition and subtractions- the result has the same number of decimal places as the measurement with the fewest decimal places. For multiplication and division- the result contains the same number of significant figures as the measurement with the fewest significant figures. 1.6 Dimensional Analysis Dimensional analysis- units are multiplied together or divided into each other along with their corresponding numerical values. The key to dimensional analysis is the correct use of conversion factors. A conversion factor is a fraction whose numerator and denominator are the same quantity expressed in different units. Ex- 12in/1ft The key of dimensional analysis is to make sure you end with your desired units or there was a mistake made. Atoms, Molecules, and Ions 12/29/15 8:28 PM 2.1 The Atomic Theory of Matter Atoms are the basic building blocks of matter and are the smallest units of an element that can combine with other elements. Dalton- law of constant composition Law of multiple proportions- if two elements A and B combine to form more than one compound, the masses of B that can combine with a given mass of A are in the ratio of small whole numbers. Dalton’s Atomic Theory: 1) Each element is composed of extremely small particles called atoms. 2) All atoms of a given element are identical, but the atoms of one element are different from the atoms of all other elements. 3) Atoms of one element cannot be changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions. 4) Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms. 2.2 The Discovery of Atomic Structure The atom is composed of subatomic particles. Particles with the same charge repel one another, whereas particles with unlike charges attract one another. Cathode rays originated at the negative electrode and traveled to the positive electrode. Rays could not be seen but their presence was detected because they cause certain materials to fluoresce, or to give off life. Thomson described cathode rays as streams of negatively charged particles, his paper became accepted as the discovery of what became known as the electron. Electrons move from negative cathodes to positive anodes. The Cathode-ray tube made it possible to calculate 1.76x10^8 coulombs per gram for the ratio of the electron’s electrical charge to its mass. Electron mass: 9.10x10^-28 g Spontaneous emission of radiation is called radioactivity. Thomson reasoned that electrons contribute only a very small fraction of an atom’s mass, they probably are responsible for an equally small fraction of an atom’s size. Thomson also proposed that the atom consists of a uniform positive sphere of matter in which the mass is evenly distributed and in which electrons are embedded like raisins in a pudding or seeds in a watermelon- this is known as the plum-pudding model. Rutherford- gold foil experiment: Rutherford was studying the angels at which alpha particles were deflected, or scattered, as they passed through a thin sheet of gold foil. Almost all of the particles passed through. Rutherford explained his results with the nuclear model of the atom, where most of the mass of each gold atom and all of its positive charge reside in a very small, extremely dense region called the nucleus. Rutherford also claimed that most of the volume is empty space where electrons move around the nucleus. Subatomic particles: protons, neutrons and electrons. The protons are positive particles in the nucleus. Neutrons are neutral particles in the nucleus. 2.3 The Modern View of Atomic Structure Every atom has an equal number of electrons and protons, so atoms have no net electrical charge. The charge of an electron is 1- and its mass is 5.486x10^-4 amu The charge of a proton is 1+ and its mass is 1.0073 amu The charge of a neutron is 0 and its mass is 1.0087 amu (amu)- atomic mass unit Atomic number- the number of protons in a particular element. Atomic number = number of protons in an element. Atoms of a given element can differ in the number of neutrons they contain. Mass number is the (atomic number) protons and neutrons added together. ty Atoms with the same atomic number but different mass number are called isotopes of one another. 2.4 Atomic Weights Atoms of different elements have different masses. Atomic mass unit can be derived from C12 assigning it a mass of exactly 12amu. The element’s atomic weight is the average atomic mass of an element by taking the sum of the masses of the isotopes multiplied by their relative abundance. 2.5 The Periodic Table Elements are arranged in order of increasing atomic number, with elements having similar properties placed in vertical columns. Periods are the horizontal rows. The vertical columns are groups. Elements in a group exhibit similarities in physical and chemical properties. Group 1A- alkali metals Group 2A- alkaline earth metals Group 6A- chalogens Group 7A- halogens Group 8A- noble gases Metals are separated from nonmetals by a stepped line, elements known as metalloids. 2.6 Molecules and Molecular Compounds Molecular form- two or more of the same type of atom bound together found in nature. A molecule made up of two atoms is a diatomic molecule. Compounds composed of molecules that contain more than one type of atom are called molecular compounds. Molecular formula-a chemical formula that indicates the actual numbers of atoms in a molecule. Empirical formula- chemical formulas that only give the relative number of atoms of each type in a molecule. Structural formula- shows which atoms are attached to which. Atoms are represented by their chemical symbols and lines are used to represent the bonds that hold the atoms together. Ball-and-stick models show atoms as spheres and bonds as sticks. Space-filling model depicts what a molecule would look like if the atoms were scaled up in a size. 2.7 Ions and Ionic Compounds Some atoms can readily gain or lose electrons. If electrons are removed from or added to an atom it becomes a charged particle called an ion. Positively charged ions are cations- they lose electrons gaining a + charge. Negatively charged ions are anions- they gain electrons gaining a – charge. In general, metals tend to lose electrons to form cations and nonmetals tend to gain electrons to form anions. Thus, ionic compounds tend to be composed of metals bonded with nonmetals. Ex- NaCl Polyatomic ions are atoms joined together in a molecule that carry a net positive or negative charge. The addition or removal of an electron produces a charged species with behavior with very different from that of its associated atom or group of atoms. Group 8A- noble gases: chemically nonreactive elements that form very few compounds. Many atoms that gain or lose electrons tend to lose electrons to end up with the same number of electrons as the noble gas closest to them on the periodic table. Group 1A- form 1+ ions Group 2A form 2+ ions Group 6A form 2- ions Group 7A form 1- ions An ionic compound is made up of cations and anions. Molecular compounds are generally composed of nonmetals only. Ex- H2O If the charges on the ions in an ionic compound are not equal, the charge on one ion (without the sign) will become the subscript on the other ion. 2.8 Naming Inorganic Compounds Chemical nomenclature- the system used in naming substances. Cations: -Cations formed from metal atoms have the same name as the metal. Ex- Na+ is sodium ion -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+ is iron(II) ion Fe3+ is iron (III) ion Ions of the same element that have different charges have different properties. Ex- colors. Most metals that form cations with different charges are transition metals, elements that occur in the middle of the periodic table. -The suffix –ous can be used to represent the ion with the lower charge and –ic represents the ion with the higher charge. Ex- Fe2+ (ferrous ion) Fe3+ (ferric ion) -Cations formed from nonmetal atoms have the names that end in –ium. Ex- H3O+ hydronium ion Anions: -The names of monoatomic anions are formed by replacing the ending of the name of the element with –ide. Ex- H- is a hydride ion. -Polyatomic anions containing oxygen have names ending in either –ate or - ite and are called oxyanions. The –ate is used for the most common or representative oxyanion of an element and –ite is used for the oxyanion that has the same charge but one fewer O atom. Ex- NO3- nitrate ion NO2- nitrite ion -Prefixes are used when the series of oxyanions of an element extends to four members, as with the halogens. The prefix per- indicates one or more O atom than the oxyanion ending in –ate. The prefix hypo- indicates one O atom fewer than the oxyanion ending in –ite. Ex- ClO4- perchlorate ion (one more O atom than chlorate) ClO3- chlorate ion ClO2- chlorite atom (one O atom fewer than chlorate) ClO- hypochlorite ion (one O atom fewer than chlorite) -Anions derived by adding H+ to an oxyanion are named by adding as a prefix the word hydrogen or dihydrogen as appropriate: Ex-CO3^2- carbonate ion HCO3^- hydrogen carbonate ion Each H+ added reduces the negative charge of the parent anion by one. The prefix hydrogen can be substituted by adding bi- in front of the preexisting ion. Ex- HCO3^- can be called bicarbonate ion Ionic compounds: -Names of ionic compounds consist of the cation name followed by the anion name. Ex- CaCl2 is calcium chloride Cu(ClO4)2 is copper (II) perchlorate An acid is a substance whose molecules yield hydrogen ions (H+) when dissolved in water. Acids: -Acids containing anions whose names end in –ide are named by changing the –ide ending to –ic and adding the prefix hydro- to the anion name and then following with the word acid. Ex- Cl- is chloride and HCl is hydrochloric acid -Acids containing anions whose names end in –ate or –ite are named by changing –ate to –ic and –ite to –ous then adding the word acid. Ex- ClO4- perchlorate ion à HClO4 perchloric acid ClO3- chlorate ion à HClO3 chloric acid ClO2- chlorite atom à HClO2 chlorous acid ClO- hypochlorite ion à HClO hypochlorous acid The procedures for naming binary (two-element) molecular compounds: -The name if the element farther to the left in the periodic table is usually written first. Exception to the rule is when the compound contains oxygen and chlorine, bromine, or iodine (any halogen except fluorine), in which case oxygen is written last. -If both elements are in the same group, the one closer to the bottom of the table is named first. -The name of the second element is given an –ide ending. -Greek prefixes indicate the number of atoms of each element. 2.9 Some Simple Organic Compounds Organic chemistry- the study of compounds of carbon. Compounds that contain carbon and hydrogen, often in combination with oxygen, nitrogen, or other elements are called organic compounds. Hydrocarbons- compounds that contain only carbon and hydrogen. Alkanes are the simplest class of hydrocarbons in which each carbon is bonded to four other atoms. Alkanes end in –ane. Alkanes with one C atom have the prefix meth- Alkanes with two C atoms have the prefix eth- Alkanes with three C atoms have the prefix prop- Alkanes with five or more carbons take their prefixes from the Greek prefixes ending in –ane. Ex- octane is C8H18 Other classes of organic compounds are obtained by replacing one or more hydrogen atoms in and alkane and adding in a functional group. Isomers- compounds with the same molecular formula but different arrangements of atoms. Structural isomers are compounds having the same molecular formula but different structural formulas. Organic compounds can form long chains of carbon-carbon bonds. Chemical Reactions and Reaction Stoichiometry 12/29/15 8:28 PM Stoichiometry- the area of study that examines the quantities of substances consumed and produced in chemical reactions. Stoichiometry is built on an understanding of atomic masses, chemical formulas and the law of conservation of mass. 3.1 Chemical Equations Reactants are the starting substances to the left of the arrow. Products are to the right of the arrow and represent the substances produced in the reaction. Chemical equations must be balanced, having an equal number of each element on each side of the arrow, because atoms are neither created nor destroyed in any reaction. Never change subscripts when balancing equations, instead place a coefficient in front of the formula since it will change the amount of the substance and not the substance’s identity. Symbols indicating the physical state of each reacant and product are often shown in chemical equations using the symbols (g) for gases, (l) for liquids, (s) for solids and (aq) for substances dissolved in aqueous (water) solutions. 3.2 Simple Patterns of Chemical Reactivity Combination reactions- two or more substances react to form one product. Ex- 2Mg (s) + O2 (g) à 2MgO (s) A combination reaction between a metal and a nonmetal produces an ionic solid. Decomposition reaction- one substance undergoes a reaction to produce two or more other substances. Combustion reactions- rapid reactions that produce a flame. Most combustion reactions we observe involve O2 from air as a reactant. Combustion of oxygen-containing derivatives of hydrocarbons produces CO2 and H2O. The rule that hydrocarbons and their oxygen-containing derivatives form CO2 and H2O when they burn in air summarizes the reactions of about 3 million compounds with oxygen. Reactions that involve intermediate steps are described as oxidation reactions instead of combustion reactions. 3.3 Formula Weights Chemical formulas and chemical equations both have a quantitative significance in that the subscripts in formulas and the coefficients in equations represent precise quantities. To calculate amounts of reactants needed to obtain a given amount of product, or otherwise extrapolate quantitative information from a chemical equation or formula, we need to know more about the masses of atoms and molecules. Formula weight (FW) of a substance is the sum of the atomic weights (AW) of the atoms in the chemical formula of the substance. If the chemical formula is that of a molecule, the formula weight is also called the molecular weight (MW). The formula weight of ionic substances is based of the sum of the atomic weights of the atoms in the empirical formula. Percentage composition of a compound is the percentage by mass contributed by each element in the substance. Elemental composition is the percentage composition of any element in a substance. where the mass of X is the number of atoms of the element x the atomic weight of the element. 3.4 Avogadro’s Number and the Mole The mole, in chemistry, is the counting unit for numbers of atoms, ions, or molecules in a laboratory-size sample. One mole is the amount of matter that contains as many objects (atoms, molecules, or whatever other objects we are considering) as the number of atoms in exactly 12g of isotopically pure C12. A mole is 6.02x10^23, which is known as Avogadro’s number. A mole is always the same number (6.02x10^23), but 1-mol samples of different substances have different masses. The atomic weight of an element in atomic mass units is numerically equal to the mass in grams of 1 mol of that element. The same numerical relationship exists between formula weight (amu) and mass (g) of 1 mol of a substance. The mass in grams of one mole of a substance is called the molar mass (g/mol) of the substance. To convert from grams to moles: Moles wanted= moles given / molar mas of chemical formula Grams wanted= moles given x molar mass of chemical formula To convert between masses and number of particles: Molecules wanted=(grams given / grams of molecule ) x 6.02x10^23 Atoms wanted of single element= molecules of sample x atoms of element 3.5 Empirical Formulas from Analyses The ratio of the number of moles of all elements in a compound gives the subscripts in the compound’s empirical formula. Using atomic weights to get molar masses, we can calculate the number of moles of each element in the sample. Then we divide the larger number by the smaller number to obtain the mole ratio. Percentages given for grams / grams of element = moles Divide each number of moles by the smallest number of moles Assign empirical value to each element. To obtain the molecular formula for any compound from its molecular formula: Whole-number multiple= molecular weight/empirical formula weight Multiply every element’s subscript in empirical formula by whole-number multiple obtained. Combustion analysis- one technique used for determining empirical formulas in the laboratory, commonly containing principally carbon and hydrogen. Example- 3.6 Quantitative Information from Balanced Equations The coefficients in a balanced chemical equation indicate both the relative number of molecules(formula units) in the reaction and the relative numbers of moles. The procedure for calculating amounts of reactants consumed or products formed in a reaction: Example- 3.7 Limiting Reactants In chemical reactions when one reactant is used up before the others, the reaction stops as soon as any reactant is totally consumed leaving excess reactants left over. Limiting reactant- the reactant that is completely consumed in a reaction because it determines, or limits, the amount of product formed. The other reactants are sometimes called excess reactants. Theoretical yield- the quantity of the product calculated to form when all of a limiting reactant is consumed. Actual yield- the amount of the product actually obtained. Percent yield- related the actual and theoretical yields of a reaction. Reactions in Aqueous Solutions 12/29/15 8:28 PM 4.1 General Properties of Aqueous Solutions A solution is a homogenous mixture of two or more substances. A solution in which water is the dissolving medium is called an aqueous solution. The substance present in the greatest quantity is usually the solvent. The other substances are the solutes and are said to be dissolved in the solvent. An electrolyte is a substance in whose aqueous solutions contain ions. Ex- NaCl A nonelectrolyte is a substance that does not form ions in solutions. Ex- C12H22O11 Ionic solids dissociate into its component ions as it dissolves. As an ionic compound dissolves, the ions become surrounded by H2O molecules and the ions are said to be solvated. Solvated ions are denoted as (aq). Solvation helps stabilize the ions in solution and prevents cations and anions from recombining, since the ions and their shells of surrounding water molecules are free to move about, the ions become uniformly dispersed throughout the solution. Strong electrolytes are those solutes that exist in solution completely or nearly completely as ions. Essentially all water-soluble ionic compounds (NaCl) and a few molecular compounds (HCl) are strong electrolytes. Weak electrolytes are those solutes that exist in solution mostly in the form of neutral molecules with only a small fraction in the form of ions. Ex- CH3COOH, acetic acid Half arrows pointing in opposite directions mean the reaction is significant in both directions. The balance between these opposing processes determines the relative numbers of ions and neutral molecules. This balance produces a state of chemical equilibrium. Chemical equilibrium- a state at which the relative numbers of each type of ion or molecule in the reaction are constant over time. A single reaction arrow is used for reactions that largely go forward, such as the ionization of strong electrolytes. Water-soluble ionic compounds are strong electrolytes. 4.2 Precipitation Reactions Precipitate reactions- reactions that result in the formation of an insoluble product. A precipitate is an insoluble solid formed by a reaction in solution. Precipitation reactions occur when pairs of oppositely charged ions attract each other so strongly that they form an insoluble ionic solid. The solubility of a substance at a given temperature is the amount of the substance that can be dissolved in a given quantity of solvent at the given temperature. Any substance with a solubility less than 0.01 mol/L is considered insoluble. To predict whether a precipitate forms when we mix aqueous solutions of two strong electrolytes we must: 1) note the ions present in the reactants 2) consider the possible cation-anion combinations 3) use solubility table to determine if any of these combinations are insoluble Reactions in which cations and anions appear to exchange partners conform to the general equation: AX +BY à AY + BX These reactions are exchange reactions or metathesis reactions. To complete and balance the equation for a metathesis reaction, we follow these steps: 1)Use the chemical formulas of the reactants to determine which ions are present. 2) Write the chemical formulas of the products by combining the cation from one reactant with the anion of the other, using the ionic charges to determine the subscripts in the chemical formulas. 3)Check the water solubilities of the products. For a precipitation reaction to occur, at least one product must be insoluble in water. 4)Balance the equation. A molecular equation shows the complete chemical formulas of reactants and products without indicating ionic character. A complete ionic equation is written with all soluble strong electrolytes shown. Spectator ions- ions that appear in identical forms on both sides of a complete ionic equation and play no direct role in the reaction. Net ionic equation- an equation that includes only the ions and molecules directly involved in the reaction, spectator ions are omitted. The sum of the ionic charges must be the same on both sides of a balanced net ionic equation because charge is conserved in reactions. If every ion in a complete ionic equation is a spectator, no reaction occurs. The net ionic equation demonstrates that more than one set of reactants can lead to the same net reaction. The complete ionic equation identifies the actual reactants that participate in a reaction. To write a net ionic equation: 1)Write a balanced molecular equation for the reaction. 2)Rewrite the equation to show the ions that form in solution when each soluble strong electrolyte dissociates into its ions. Only strong electrolytes dissolved in aqueous solution are written in ionic form. 3)Identify and cancel spectator ions. Example- 4.3 Acids, Bases, and Neutralization Reactions Acids and bases are common electrolytes. Acids are substances that ionize in aqueous solution to form hydrogen ions H+. Acids are proton donors. Protons in aqueous solution are solvated by water molecules, just as other cations. Monoprotic acids yield one H+ per molecule of acid. Diprotic acid yields two H+ per molecule of acid. Bases are substances that accept H+ ions. Bases produce hydroxide ions (OH-) when they dissolve in water. Most common bases: NaOH, KOH, and Ca(OH)2. Compounds that do not conatin OH- ions can also be bases. Ex- ammonia NH3 accepts a H+ ion from a water molecule and thereby produces an OH- ion. Acids and bases that are strong electrolytes (completely ionized in solution) are strong acids and strong bases. Those that are weak electrolytes (partly ionized) are weak acids and weak bases. How to identify strong and weak electrolytes: When a solution of an acid and a solution of a base are mixed, a neutralization reaction occurs. The products of the reaction have none of the characteristic properties of either the acidic solution or the basic solution. A neutralization reaction between an acid and a metal hydroxide produces water and a salt. A salt has come to mean any ionic compound whose cation comes from a base and whose anion comes from an acid. Neutralization reactions between acids and metal hydroxides are metathesis reactions. Many bases besides OH- react with H+ to form molecular compounds. Two of these that you might encounter in the laboratory are the sulfide ion and the carbonate ion. Both of these anions react with acids to form gases that have low solubilities in water. Carbonates and bicarbonates react with acids to form CO2(g). Example- 4.4 Oxidation-Reduction Reactions Oxidation-reduction reactions or redox reactions are a kind of reaction in which electrons are transferred from one reactant to another. The most familiar redox reactions is corrosion of a metal. Corrosion is the conversion of a metal into a metal compound, by a reaction between the metal and some substance in its environment. When a metal corrodes, each metal looses electrons and forms a cation, which can form with an anion to form an ionic compound. When an atom, ion, or molecule becomes more positively charged (loses electrons), it is oxidized. Loss of electrons by a substance is called oxidation. Many metals react directly with O2 in air to form metal oxides. In these reactions the metal loses electrons to oxygen, forming an ionic compound of the metal ion and oxide ion. When an atom, ion, or molecule becomes more negatively charged (gains electrons) we say that it is reduced. The gain of electrons by a substance is called reduction. Each atom in a neutral substance or ion is assigned an oxidation number. For monoatomic ions- the oxidation number is the same as the charge. For neutral molecules and polyatomic ions, the oxidation number of a given atom is a hypothetical charge. This charge is assigned by artificially dividing up the electrons among the atoms in the molecule or ion. We use the following rules for assigning oxidation numbers: 1) For an atom in its elemental form- tne oxidation number is always zero. 2) For any monatomic ion- the oxidation number equals the ionic charge. Ex- K+ has an oxidation number of +1 3)Nonmetals usually have negative oxidation numbers, although they can sometimes be positive: A) The oxidation number of oxygen is usually 2- in both ionic and molecular compounds. The major exception is in compounds called peroxides, which contain the O2^2- ion, giving each oxygen an oxidation number of -1. B) The oxidation number of hydrogen is usually +1 when bonded to nonmetals and -1 when bonded to metals. C) The oxidation number of fluorine is -1 in all compounds. The other halogens have an oxidation number of -1 in most binary compounds. When combined with oxygen, as in oxyanions, however, they have positive oxidation states. 4) The sum of the oxidation numbers of all atoms in a neutral compound is zero. The sum of the oxidation numbers in a polyatomic ion equals the charge of the ion. Ex- H3O+ contains H+ (aq) and O-2 so the sum of the oxidation numbers is 3(+1) + (-2)= +1 The reaction between a metal and either an acid or a metal salt conforms to the general pattern: A + BX à AX + B Displacement reactions are reactions in which the ion in solution is displaced (replaced) through oxidation of an element. Many metals undergo displacement reactions with acids, producing salts and hydrogen gas. Whenever one substance is oxidized, another substance must be reduced. A list of metals arranged in order of decreasing ease of oxidation is called an activity series. The metals at the top of the table (the alkali metals and the alkaline earth metals) are most easily oxidized because they react most readily to form compounds and are called active metals. Metals at the bottom of the activity series, such as the transition elements from Group 8B and 1B, are very stable and form compounds less readily. These metals are called noble metals because of their low reactivity. Any metal on the list can be oxidized by the ions of elements below it. For example- Copper is above silver in the series which means copper metal is oxidized by silver ions: Cu(s) + 2Ag+(aq) à Cu2+(aq) + 2Ag(s) The oxidation of copper to copper ions is accompanied by the reduction of silver ions to silver metals. Only metals above hydrogen in the activity series are able to react with acids to form H2. 4.5 Concentrations of Solutions Concentration is the amount of solute dissolved in a given quantity of solvent or quantity of solution. Molarity (M) expresses the concentration of a solution as the number of moles of solute in a liter of solution: To find the moles given molarity: Moles wanted= volume given x (moles given/ 1 L of solution) Stock solutions- solutions used routinely in the laboratory often purchased or prepared in concentrated form. Dilution- adding water to solutions of lower concentrations. Moles of solute before dilution= moles of solute after dilution For dilution problems: Find the moles of diluted solution Take (moles of diluted solution) x (1 L of solution/ moles of diluted solution)= Liters of concentration of the solution The number of moles of solute is the same in both the concentrated and dilute solutions and that moles = molarity x liters: Moles of solute in concentrated solution = moles of solute in diluted solution Mconc x Vconc = Mdil x Vdil 4.6 Solution Stoichiometry and Chemical Analysis Titration is a process used to carry out the concentration of a particular solute in a solution. Titration involves combining a solution where the solute concentration is not known with a reagent solution of known concentration, otherwise called a standard solution. Equivalence point is the point at which stoichiometrically equivalent quantities are brought together. Thermochemistry 12/29/15 8:28 PM Thermodynamics- the study of energy and it’s transformations. Thermochemistry- the relationship between chemical reactions and energy changes that involve heat. 5.1 Energy Energy- the capacity to do work or transfer heat. Work- the energy used to cause an object to move against a force. Heat- the energy used to cause the temperature of an object to increase. Kinetic energy- the energy of motion. The kinetic energy of an object increases as its speed increases. An object has potential energy by virtue of its position relative to other objects. Potential energy is, in essence, the “stored” energy that arises from the attractions and repulsions an object experiences in relation to other objects. Gravitational forces play a negligible role in the ways that atoms and molecules interact with one another. One of the most important forms of potential energy in chemistry is electrostatic potential energy, Eel. This arises from the interactions between charged particles. SI unit for energy is the joule (J) A calorie (cal) was originally defined as the amount of energy required to raise the temperature of 1g of water from 14.5 to 15.5°C. A calorie is now defined in terms of the joule: 1 cal=4.184 J The portion we single out for study is called the system. Everything else is called the surroundings. Systems may be open, closed, or isolated. An open system is one in which matter and energy can be exchanged with the surroundings. A closed system is a system that can exchange energy but not matter with their surroundings. An isolated system is one in which neither energy nor matter can be exchanged with the surroundings. Work is the energy transferred when a force moves an object. A force is any push or pull exerted on an object. W=F x d Heat is the energy transferred from a hotter object to a colder one. 5.2 The First Law of Thermodynamics The first law of thermodynamics is that energy can be neither created or destroyed. Any energy that is lost by a system must be gained by the surroundings, and vice versa.   Internal energy (E) of a system is the sum of all the kinetic and potential energies of the components of the system. The change in internal energy, denoted ΔE, as the difference between Efinal and Einitial. Change in internal energy has three parts: 1) A number 2) A unit which gives magnitude of the change 3) A sign that gives the direction A positive delta E means the system has gained energy from its surroundings (final is greater than initial) A negative delta E value indicates the system has lost energy to its surroundings (initial is greater than final). In a chemical reaction the initial state is the reactants and the final state is the products. Example- Endothermic- a process in which the system absorbs heat. Heat flows into the system from its surroundings. Exothermic- a process in which the system loses heat. Heat exits or flows out of the system into the surroundings. State function- a property of a system that is determined by specifying the system’s condition, or state (in terms of temperature, pressure, and so forth). The value of a state function depends only on the present state of the system, not on the path the system took to reach that state. Ex- enthalpy, entropy, internal energy, pressure, volume 5.3 Enthalpy Enthalpy (H)- the internal energy plus the product of pressure (P) and volume (V) of the system. The work involved in the expansion or compression of gases is called pressure-volume work (P-V work). When pressure is constant in a process, the sign and magnitude of the pressure-volume work are given by: w = -P  ΔV Where P is pressure and ΔV = final volume – initial volume, which is the change in volume of the system. The pressure P is always either a positive number or zero. Enthalpy equals the change in internal energy plus the product of the constant pressure and the change in volume. ΔE=q+w And w=-P  ΔV (at constant pressure) The change in enthalpy equals the heat q pained or lost at a constant pressure. When H is positive, the system has gained heat from the surroundings, the process is endothermic. When H is negative, the system has lost heat from the surroundings, the process is exothermic. 5.4 Enthalpies of Reaction The enthalpy change for a chemical reaction is given by: A negative sign for H means the reaction is exothermic. A positive sign for H means the reaction is endothermic. Balanced chemical equations that show the associated enthalpy change in this way are called thermochemical equations. The following guidelines are helpful when using thermochemical equations and enthalpy diagrams: 1) Enthalpy is an extensive property. The magnitude of H is proportional to the amount of reactant consumed in the process. 2) The enthalpy change for a reaction is equal in magnitude, but opposite in sign to H for the reverse reaction. When we reverse a reaction, we reverse the roles of the products and the reactants. 3) The enthalpy change for a reaction depends on the states of the reactants and the products. 5.5 Calorimetry The measurement of heat flow is calorimetry. A device used to measure heat flow is a calorimeter. The temperature change experienced by an object when it absorbs a certain amount of heat is determined by its heat capacity C. The heat capacity of an object is the amount of heat required to raise its temperature by 1K or 1 degree Celcius. The greater the heat capacity, the greater the heat required to produce a given increase in temperature. The heat capacity of one mole of a substance is called its molar heat capacity, Cm. The heat capacity of one gram of a substance is called its specific heat capacity, or specific heat, Cs. The equation for specific heat is: or T in K= T in C ° Cs = 4.18 J/g-K The units for specific heat is J/g-°C We can calculate the quantity of heat a substance gains or loses by using its specific heat together with its measured mass and temperature change. Cm= (q/grams given) x (grams of substance/ per mole) heat of the solution = -heat of the reaction qsoln=-qrxn A temperature increase ( T >0) means the reaction is exothermic (qrxn <0). A combustion reaction is a reaction in which a compound reacts completely with excess oxygen. Combustion reactions are most accurately studied using a bomb calorimeter. 5.6 Hess’s Law Hess’s Law states that if a reaction is carried out in a series of individual steps, H for the overall reaction equals the sum of the enthalpy canges for the individual steps. H is a state function, so for a particular set of reactants and products, H is the same whether the reaction takes place in one step or in a series of steps. 5.7 Enthalpies of Formation Enthalpy of formation, or heat of formation, Hf, the subscript f indicates that the substance has been formed from its constituent elements. The standard state of a substance is its pure form at atmospheric pressure (1 atm) and the temperature of interest, which we usually choose to be 298 K (25 degrees C). The standard enthalpy of formation of a reaction is defined as the enthalpy change when all reactants and products are in their standard states. We denote enthalpy change as H° , where the subscript ° indicates standard-state conditions. The standard enthalpy of formation of a compound , Hf° is the change in enthalpy for the reaction that forms one mole of the compound from its elements with all substances in their standard states. The standard enthalpy of formation of the most stable form of any element is zero because there is no formation reaction needed when the element is already in its standard state. 5.8 Foods and Fuels Most chemical reactions used for the production of heat are combustion reactions. The energy released when one gram of any substance is combusted is the fuel value of the substance. The fuel value of any food or fuel can be measured by calorimetry. Fossil fuels- the world’s major sources of energy. Natural gas- consists of gaseous hydrocarbons, compounds of hydrogen and carbon. Petroleum- a liquid composed of hundreds of compounds, most of which are hydrocarbons, with the remainder being chiefly organic compounds containing sulfur, nitrogen, or oxygen. Coal, which is a solid, contains hydrocarbons of high molecular weight as well as compounds containing sulfur, oxygen, or nitrogen. Renewable energy sources- sources that are essentially inexhaustible. Electronic Structure of Atoms 12/29/15 8:28 PM The electron structure of an atom refers to the number of electrons in the atom as well as their distribution around the nucleus and their energies. 6.1 The Wave Nature of Light Periodic means the pattern of peaks and troughs repeat itself at regular intervals. Wave length- the distance between two adjacent peaks or troughs. The number of complete wavelengths, or cycles, that pass a given point each second is the frequency of the wave. All electromagnetic radiation moves at the same speed, the speed of light. For a wave to have high frequency, it must have a short wavelength. Where lambda (λ) is wavelength Nu (v) is frequency C is the speed of light Various types of electromagnetic radiation arranged in order of increasing wavelength displayed is called the electromagnetic spectrum. Visible light corresponds to the wavelengths from about 400 to 750 nm. Frequency is expressed in cycles per second, a unit also called hertz (Hz). Example calculating frequency from wavelength: 6.2 Quantized Energy and Photons Plank gave his proposal of energy being either released or absorbed by atoms only in a discrete “chunks” of some minimum size the name quantum. Quantum (meaning “fixed amount”) to the smallest quantity of energy that can be emitted or absorbed as electromagnetic radiation. He proposed the energy, E, of a single quantum equals a constant times the frequency of the radiation: The h is plank’s constant. h= 6.626x10^-34 J-s Energy can be released only in specific amounts, so the energies are quantized and their values are restricted to certain quantities. Einstein used Planck’s theory to explain the photoelectric effect. The photoelectric effect states light shinning on a clean metal surface causes electrons to be emitted from the surface. A minimum frequency of light, different for different metals, is required for the emission of electrons. Radiant energy striking the metal surface behaves like a stream of tiny energy packets, each packet is like a “particle” of energy and therefore a photon. Einstein deduced that each photon must have an energy equal to Planck’s constant times the frequency of light: Example- 6.3 Line Spectra and the Bohr Model Radiation composed of a single wavelength is monochromatic. Radiation composed of many different wavelengths is polychromatic. A spectrum is produced when radiation from such sources is separated into its component wavelengths. Continuous spectrum is a rainbow of colors, containing light of all wavelengths. A spectrum containing radiation of only specific wavelengths is called a line spectrum. Rydberg equation allows us to calculate the wavelengths of all the spectral lines of hydrogen: Lambda is the wavelength of a certain spectral line Rh is the Rydberg constant – 1.096776 x 10^7 m^-1 n1 and n2 are positive integers Bohr’s three postulates: 1. Only orbits of certain radii, corresponding to certain specific energies, are permitted for the electron in a hydrogen atom. 2. An electron in a permitted orbit in an “allowed” energy state. An electron in an allowed energy state does not radiate energy and, therefore, does not spiral into the nucleus. 3. Energy is emitted or absorbed by the electron only as the electron changes from one allowed energy state to another. This energy is emitted or absorbed as a photon that has energy E=hv. Equation for the energies corresponding to the allowed orbits for the electron in the hydrogen atom: h ,c, and Rh are the Planck constant, the speed of light and the Rydberg constant . The integer n, which can have whole-number values of 1,2,3… to infinity is called the principal quantum number. Each allowed orbit corresponds to a different value of n. The radius of the orbit gets larger as n increases. The lowest energy state is called the ground state. When an electron is in a higher-energy state, the atom is said to be in an excited state. The state in which the electron is completely separated from the nucleus is called the reference, or zero-energy, state of the hydrogen atom. Equation for energies jumping energy states: E is positive when the nf is greater than ni because the electron is jumping to a higher-energy orbit. E is negative when nf is less than ni because the electron is falling in energy to a lower-energy orbit. The energy of the photon (Ephoton) must equal the difference in energy between the two states ( E). When E is positive, a photon must be absorbed as the electron jumps to a higher energy. When E is negative, a photon is emitted as the electron falls to a lower energy level. We recognize that the line spectra are the results of emission, so for these transitions. The existence of discrete spectral lines can be attributed to the quantized jumps of electrons between energy levels. The equation used to solve that is: What is most specific about Bohr’s model is that it introduces two important ideas that are also incorporated into our current model: 1.Electrons exist only in certain discrete energy levels, which are described by quantum numbers. 2.Energy is involved in the transition of an electron from one level to another. 6.4 The Wave Behavior of Matter De Broglie suggested that an electron moving about the nucleus of an atom behaves like a wave and therefore has a wavelength. He proposed that the wavelength of the electron, or of any other particle, depends on its mass, m, and its velocity, v, where h is Planck’s constant. The quantity mv for any object is called its momentum. De Broglie used the term matter waves to describe the wave characteristics of material particles. Example: X-ray diffraction: a phenomenon when electrons pass through a crystal, an interference pattern results that is characteristic of wave-like properties of electromagnetic radiation. As electrons pass through a crystal, they are similarly diffracted. Thus, a stream of moving electrons exhibits the same kinds of wave behavior as X rays and all other types of electromagnetic radiation. Heisenberg’s uncertainty principle states that it is impossible for us to know simultaneously both the exact momentum of the electron and its exact location in space. 6.5 Quantum Mechanics and Atomic Orbitals Probability density or electron density- the square of the wave function, Ψ^2, at a given point in space represents the probability that the electron will be found at the location. An orbital is a quantum mechanical-model which describes the electrons in terms of probability. Each orbital has a characteristic shape and energy. The Bohr model introduced a single quantum number, n, to describe an orbit. The quantum number mechanical model uses three quantum numbers, n, l, and ml which result naturally from the mathematics used to describe an orbital. 1. The principal quantum number, n, can have positive integral values 1,2,3… As n increases, the orbital becomes larger, and the electron spends more time further 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. 2. The second quantum number, the angular momentum quantum number, l, can have integral values from 0 to (n-1) for each value of n. This quantum number defines the shape of the orbital. 3. The magnetic quantum number, ml, can have integral values between –l and l, including 0. This quantum number describes the orientation of the orbital in space. Electron shell- the collection of orbitals with the same value of n. The set of orbitals that have the same n and l values is called a subshell. The shell with principal quantum number n consists of exactly n subshells. Each subshell consists of a specific number of orbitals corresponding to a different allowed value for ml. The total number of orbitals in a shell is n^2, where n is the principal quantum number of the shell. 6.6 Representation of Orbitals The s orbital: Radial probability density- the probability that the electron is at a specific distance from the nucleus. Comparing the radial probability distributions for the 1s, 2s, 3s orbitals reveals three trends: 1. For an ns orbital, the number of peaks is equal to n, with the outermost peak being larger than inner ones. 2. For an ns orbital, the number of nodes is equal to n-1. 3. As n increases, the electron density becomes more spread out, that is, there is a greater probability of finding the electron further from the nucleus. Example of the s orbitals: The p orbitals: The orbitals for which l=1 are the p orbitals. Each p subshell has three orbitals, corresponding to the three allowed values of ml: -1,0,1. Beginning with the n=2 shell, each shell has three p orbitals. For each value of n, the three p orbitals have the same size and shape but differ from one another in spatial orientation. P orbitals increase in size as we move from 2p, 3p, 4p, and so forth. The d and f orbitals: When n is 3 or greater, we encounter the d orbitals (for which l=2). There are five 3d orbitals, five 4d orbitals, and so forth because in each shell there are five possible values for the ml quantum number: -2,-1, 0, 1, 2. The different d orbitals in a given shell have different shapes and orientations in space. When n is 4 or greater, there are seven equivalent f orbitals (for which l=3). The shapes of f orbitals are even more complicated than those of the d orbitals and are not presented here. 6.7 Many–Electron Atoms In a many-electron atom, for a given value of n, the energy of an orbital increases with increasing value of l. Orbitals with the same energy are said to be degenerate. Qualitative energy-level diagram: the exact energies of the orbitals and their spacings differ from one atom to another. Electron spin- each electron behaves as if it were a tiny sphere spinning on its own axis. A new quantum number is spin magnetic quantum number, ms (s stands for spin). The possible values are allowed for ms, -1/2 and +1/2, which were first interpreted as indicating two opposite directions in which the electron can spin. The Pauli exclusion principle states that no two electrons in an atom can have the same set of four quantum numbers: n, l, ml, ms. In a given orbital, the values of m, l, and ml are fixed. If we want to put more than one electron in an orbital and satisfy the Pauli Exclusion Principle, our only choice is to assign different ms values to the electrons. An orbital can hold a maximum of two electrons and they must have opposite spins. 6.8 Electron Configuration Electron configuration- the way electrons are distributed among the various orbitals of an atom. The most stable electron configuration, the ground state, is that in which electrons are in the lowest possible energy states. The orbitals are filled in order of increasing energy, with no more than two electrons per orbitals. An orbital diagram- each orbital is denoted by a box and each electron by a half arrow. A half arrow pointing up ( ) represents an electron with a positive spin magnetic quantum number (ms = +1/2) and a half arrow pointing down ( ) represents an electron with a negative spin magnetic quantum number (ms = -1/2). Electrons having opposite spins are said to be paired when they are in the same orbital ( ). An unpaired electron is not accompanied by a partner of opposite spin. Hund’s rule- for degenerate orbitals, the lowest energy is attained when the number of electrons having the same spin is maximized. This means that electrons occupy orbitals singly to the maximum extent possible and that these single electrons in a given subshell all have the same spin magnetic quantum number. Hund’s rule is based on the fact that electron’s repel one another. By occupying different orbitals, the electrons remain as far as possible from one another, thus minimizing electron-electron repulsions. In writing the condensed electron configuration of an element, the electron configuration of the nearest noble-gas element or lower atomic number is represented by it’s chemical symbol in brackets. Example- Inner-shell electrons are referred to as the core electrons. The outer shell electrons include electrons involving chemical bonding, which are called the valence electrons. The transition elements or transition metals start with the fourth row on the periodic table and is ten elements and wider than the two previous rows. There are seven degenerate 4f orbitals, each corresponding to the seven allowed values of ml, ranging from -3 to 3. It takes 14 electrons to fill the seven allowed orbitals completely. The 14 elements corresponding to the filling of the 4f orbitals are known as either the lanthanide elements or the rare earth elements. The actinide elements begin by filling the 7s orbital and are radioactive, and most of them are not found in nature. 6.9 Electron Configuration and the Periodic Table The s and p block together are the representative elements or the main- group elements. 2,6,10 and 14 are the number of electrons that can fill the s,p,d, and f subshells respectively. The sum of electrons in increasing subshells in electron configuration should equal the atomic number of the element. For representative elements, we do not consider the electrons in completely filled d or f subshells to be valence electrons. For transition elements, we do not consider the electrons in a completely filled f subshell to be valence electrons. Anomalous behavior is largely a consequence when there are enough electrons to form precisely half-filled sets of degenerate orbitals (as chromium) or a completely filled d subshell (as in copper). There are a few similar cases among the heavier transition metals (those with partially filled 4d or 5d orbitals) and among the f-block metals. Periodic Properties of the Elements 12/29/15 8:28 PM 7.1 Development of the Periodic Table Dimitri Mendeleev and Lothar Meyer both noted similar chemical and physical properties recur periodically when the elements are arranged in order of increasing atomic weight. Mendeleev is given credit for advancing his ideas because of his insistence that elements with similar characteristics be listed in the same column which forced him to leave blank spaces in his table. Henry Moseley developed the concept of atomic numbers. 7.2 Effective Nuclear Charge The attractive force between an electron and the nucleus depends on the magnitude of the nuclear charge and on the average distance between the nucleus and the electron. The force increases as the nuclear charge increases and decreases as the electron moves farther from the nucleus. Electron-electron repulsions cancel some of the attraction of the electron to the nucleus so that the electron experiences less attraction than it would if the other electrons weren’t there. In essence, each electron in a many-electron atom is screened from the nucleus by the other electrons and it therefore experiences a net attraction that is less than it would in the absence if other electrons. Effective nuclear charge, Zeff is the partially screened nuclear charge. Because the full attractive force of the nucleus has been decreased by the electron repulsions, we see that the effective nuclear charge is always less than the actual nuclear charge (Zeff < Z). We can define the amount of screening of the nuclear charge quantitatively by using a screening constant, S, such that Zeff = Z- S Where Z is the atomic number Where S the number of core electrons of the element, which will always be a positive number. The effective nuclear charge increases from left to right across the periodic table. The effective nuclear charge increases slightly going down a column. 7.3 Sizes of Atoms and Ions The bonding atomic radius for any atom in a molecule is equal to half the bond distance d. The bonding atomic radius (also known as the covalent radius) is smaller than the nonbonding atomic radius. Atomic radius increases down a group. Atomic radius decreases from left to right across a period. Cations are smaller than their parent ions. Cations increase down a group and decrease from left to right across a period. Anions are larger than their parent ions. Anions increase down a group and increase from left to right across a period. For ions carrying the same charge, ionic radius increases as we move down a column on the periodic table. An isoelectronic series- a group of ions all containing the same number of electrons. In any isoelectric series we can list the members in order of increasing atomic number. 7.4 Ionization Energy The ionization energy of an atom or ion is the minimum energy required to remove an electron from the ground state of the isolated gaseous atom or ion. First ionization energy, I1, is the energy needed to remove the first electron from a neutral atom. Na (g) à Na+ (g) + e- Second ionization energy, I2, is the energy needed to remove an electron. Na+ (g) à Na2+ + e- The greater the ionization energy, the more difficult it is to remove an electron. First ionization energy generally increases as we move across a period. First ionization energy generally decreases as we move down a column. The s and p blocks show a larger range of first ionization energy values than do the transition metals. As we move across a period there is both an increase in effective nuclear charge and a decrease in atomic radius, causing the ionization energy to increase. When electrons are removed from an atom to form a cation, they are always removed first from the occupied orbitals having the largest principal quantum number, n. An example of removal of electrons from a lithium atom: 7.5 Electron Affinity All ionization energies for atoms are positive: Energy must be absorbed to remove an electron. Most atoms can also gain electrons to form anions. The energy change that occurs when an electron is added to a gaseous atom is called the electron affinity. Electron affinity measures the attraction, or affinity, of the atom for the added electrons. For most atoms, energy is released when an electron is added. Example- Electron affinities do not change greatly as we move down a group. 7.6 Metals, Nonmetals, and Metalloids The more an element exhibits the physical and chemical properties of metals, the greater its metallic character. Metallic character generally increases as we proceed down a group of the periodic table and decrease as we proceed right across a period. Metals tend to have low ionization energies and therefore tend to form cations relatively easily. As a result, metals are oxidized (lose electrons) when they undergo chemical reactions. First ionization energy is the best indicator of whether an element behaves as a metal or a nonmetal. Compounds made up of a metal


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