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Chapter 9: Chemical Bonds Compounds and Chemical Change • Chemistry is all about bonds being made or broken. • Bonds are simply the sharing or exchanging of electrons. • A chemical reaction occurs when substances are created by the breaking or formation of chemical bonds. • Example: H2 and O2 can react (bonds formed) to make H2O or H2O can decompose (bonds broken) to form H2 and O2. • Left: The nobel gases, example Neon above, are monatomic- meaning they naturally occur as stable, single atoms. Right: Many gases, such as Hydrogen, are diatomic, which means they naturally occur as a molecule with two atoms. Not shown: Some gasses, such as Ozone- a form of oxygen, occur in a three atom arrangement called triatomic. Decomposition of Water • The figure shows the decomposition of water into H2 and O2. The reaction also works in reverse to make water from H2 and O2. • Note: both H2 and O2. are diatomic gases.. Atoms and Molecules • An atom is the smallest unit of an element that can exist alone or in combination with other elements. • A molecule is the smallest particle of a compound, or a gaseous element, that can exist and still retain the characteristic chemical properties of a substance. Chemical Reaction • A chemical reaction is a change in matter in which different chemical substances are created by forming or breaking chemical bonds. – Chemical bonds are formed when atoms of elements bind together to form compounds. – Chemical bonds are broken when a compound is decomposed into simpler substances. Chemical Equations • A chemical equation is a symbolic representation of a chemical reaction. – The reactants are the substances that are changed in a reaction . They are always on the left side of a written equation. – The products are the substances that are formed in the reaction. They are always written on the right side of the equation. Valence Electrons • Valence electrons are those electrons that reside in the outer shell of an atom. – These are the electrons that are lost/gained or shared in a chemical reaction. – Inner electrons are in filled orbitals in lower energy levels. • These electrons are therefore not available for interactions with other electrons. The Octet Rule – Atoms attempt to acquire an outer orbital with eight electrons through chemical reactions. – This gives them an outer shell configuration like their nearest noble gas and therefore they become stable. – From the family number of the representative elements, you can determine the number of valence electrons, and therefore the number of electrons necessary to gain the stable configuration. Octet Rule Example Sodium is in Group IA and therefore has 1 electron in its outer shell -1 valence electron. • If sodium loses one electron if becomes Na+ and has 8 electrons in its outer shell. • It then has the electron configuration of Ne (neon) and has a filled outer shell configuration. Chemical Bonds • Introduction – A chemical bond that is an attractive force that holds the atoms in a compound together. – An Ionic bond forms when an atom transfers an electron to another atom during a chemical reaction. – A covalent bond forms when atoms share electrons in a chemical bond. – Metallic bonds form in metals. Bond Theories – Bonds that form in compounds can be described in several ways. – Molecular orbital theory describes the electrons as belonging to the whole molecule which gives the orbital its own shape, orientation, and energy levels. – Isolated atom description considers the electrons around the atoms as being isolated from the rest of the molecule. Ionic Bonds – Ionic bonding occurs when one atom transfers an electron to another atom. • The difference in electrical charge results in an electrostatic attraction between unlike electrical charges. • This occurs when a metal reacts with a nonmetal. • A crystal of sodium chloride is composed of sodium and chlorine ions held together by electrical attraction. Each positive sodium ion is surrounded by six negative chlorine ions, and each chlorine ion is surrounded by six sodium ions. This atomic arrangement gives the sodium chloride crystal a cubic structure. Salt Crystals • Left: The cubic structure of table salt crystals is easily seen when magnified about ten times. • Right: Model of a salt crystal. Heat of Formation – Energy and Electrons in Ionic Bonding – energy + Na+ Na+ + e– Cl + e- Cl- + energy – Na+ + Cl- Na+Cl- + energy • The energy that is released in steps 2 and 3 is greater that that absorbed in step one and an ionic bond is formed. – This energy is called the heat of formation. Ionic Bonding Rules • Two rules for keeping track of electrons in ionic bonding reactions. – Ions are formed when atoms gain or lose electrons to achieve a noble gas configuration. – The number of electrons that are lost must equal the number of electrons that are gained. Compounds/Formulas – Ionic Compounds and Formulas • The formula of a compound describes what elements are in the compound and in what proportions. • Compounds that are held together by ionic bonds are called ionic compounds. • The elements in Group IA and IIA tend to lose electrons for form positive ions • The elements in Group VIA and VIIA tend to gain electrons to form negative ions. Covalent Bonds – A covalent bond is a chemical bond that is formed when two atoms share a pair of electrons. – H. + H. H:H – Covalent Compounds and Formulas • Since a pair of electrons is shared in a covalent bond, the electrons move throughout the entire molecular orbital. • In the above example, both hydrogen atoms gain the electron configuration of helium. • Covalent compounds are compounds with covalent bonds. • Covalent compounds form from atoms on the right side of the periodic table (nonmetals). Multiple Bonds • Bonding pairs can also be represented by lines connecting atoms. H:H = H—H • When one pair of electrons is shared, it is called a single bond. • When two pairs of electrons are shared it is called a double bond. • When three pairs of electrons are shared it is called a triple bond. Electronegativity – Electronegativity is a measure of the relative ability of an atom to attract bonding electrons. – Elements with higher values have the greatest affinity for bonding electrons. – Differences in electronegativity can be used to predict whether a bond will be ionic or covalent. • If the absolute difference is 0.5 or less, the bond will be covalent. • If the absolute difference is 1.7 or more the bond will be ionic. Ionic Compound Names – Ionic compounds that are formed from metal ions are named by naming the metal ion (electropositive ion) first, followed by the nonmetal (electronegative ion). – The ending of the nonmetal is changed to end in -ide – When a metal can have various oxidation states the oxidation state is give by roman numerals in parenthesis after the name of the metal. Ionic Compound Formulas – Rules • The symbol for the positive element is written first, followed by the symbol of the negative element • Subscripts are used to indicate the numbers of ions needed to produce an electrically neutral compound. • Example calcium chloride – calcium is Ca2+ and chlorine is Cl– there needs to be two negative charges to balance the 2+ on the calcium – the formula is therefore CaCl2 Covalent Compound Names – Names for covalent compounds uses Greek prefixes to indicate numbers of atoms of each element • The first element in the formula is named first with a prefix indicating the number of atoms if the number is greater than one. • The stem name of the second element in the formula is named next, with a prefix used with the stem name if two elements can form more than one compound – The suffix –ide is again used. • Example – CO = carbon monoxide – CO2 = carbon dioxide Covalent Compound Formulas – The systematic name tells you the formula. – Formulas indicate how many atoms of one element combine with atoms of another element. – The number of covalent bonds that an atom can form is called its valence. – Lone pairs create the possibility of creating coordinate covalent compounds.