Chemistry Matter and Change: Chapter 7 Atoms in ionic compounds are held together by chemical bonds formed by the attraction of oppositely charged ions. Ions are formed when atoms gain or lose valence electrons to achieve a stable octet electron configuration. • • • Define a chemical bond. Describe the formation of positive and negative ions Relate ion formation to electron configuration. Ion Valence electron Octet Electron configuration Lewis-dot diagrams Electron affinity Chemical bond Ionic bond Cation Anion The force that holds two atoms together Three types Ionic bonds Metallic bonds Covalent bonds *Chap 7 *Chap 7 *Chap 8 Each valence electron is represented as a dot around the nuclear core of the element. The most stable electron configuration for an element is the nearest noble gas. ns2np6 Octet Ions gain or lose electrons to achieve noble gas configurations Cation: a positively charged ion Results when electrons are lost Group 1 loses 1 electron +1 charge Group 2 loses 2 electrons +2 charge Group 13 loses 3 electrons +3 charge Groups 3-12 usually lose 2 electrons Most have +2 charge (range from +1 to +3) Anion: negatively charged ion Formed when electrons are gained Non-metals Group 15 gains 3 electrons 3- charge Group 16 gains 2 electrons 2- charge Group 17 gains 1 electron 1- charge A chemical bond is the force that holds two atoms together Some atoms gain or lose electrons to gain a stable configuration; these are called ions Most stable configurations end: ns2np6. • • • Define a chemical bond. Describe the formation of positive and negative ions Relate ion formation to electron configuration. Oppositely charged ions attract each other forming electrically neutral ionic compounds. Describe the formation of ionic bonds and the structure of ionic compounds Generalize about the strength of ionic bonds based on the physical properties of ionic bonds Categorize ionic bond formation as exothermic or endothermic Compound Chemical bond Physical property Chemical property Electronegativity Ionic bond Ionic compound Crystal lattice Binary compound Electrolyte Electrons are exchanged between atoms Increases stability of both Ions are charges held together by the opposite Bond formed between two elements with an electronegativity difference > 1.7 Crystallize as sharply defined particles Formed from a metal and a non-metal Contain only two elements Examoples NaCl MgO CaCl2 Fe2O3 Net charge on all ions in a compound must be zero (0)! More on this later!!! Crystal Lattice: Highly organized crystal of cations and anions Anion Cation Crystalline shape depends on involved the ions Physical properties Very strong Solid at normal temperatures Very high melting point and boiling point Many have brilliant colors due to transition metals Hard, rigid Brittle Conductivity (ability for electric charge to move through a substance Solids have electrons locked in place Non conductive Aqueous solutions have easily moveable electrons Electrolytes Good conductors Dissolve in water May have radically different properties than the elements that compose them Polar Dissolution Formation of lattice is always exothermic. Ionic compounds contain ionic bonds formed by the attraction of oppositely charged ions. Ions in an ionic compound are arranged in a repeating pattern called a lattice. Ionic compounds are electrolytes; they conduct electricity in liquid and aqueous states. Describe the formation of ionic bonds and the structure of ionic compounds Generalize about the strength of ionic bonds based on the physical properties of ionic bonds Categorize ionic bond formation as exothermic or endothermic In written names and formulas for ionic compounds, the cation appears first, followed by the anion. • • • Relate a formula unit of an ionic compound to its composition Write formulas for ionic compounds and oxyanions. Apply naming conventions to ionic compounds and oxyanions. Anion Cation Metal Non-metal Formula unit Monatomic ion Polyatomic ion Oxidation number Oxyanion Formula unit- simplest way to indicate the composition of an ionic substance NaCl MgCl2 Cl- Mg2+ Cl- Ions in which only one element is present • Na+, Cl-, Mg2+, P3- Fancy word for “charge” aka oxidation state Transition metals may have multiple oxidation states Must tell the oxidation state Ex: Iron 2+ is Iron II; Iron 3+ is Iron III CxAy C is cation A is anion x number of cations in one unit y is number of anions in one unit CxAy Cation is always first Anion is always second Net oxidation MUST BE ZERO 1. 2. 3. 4. Write out each ion. Place oxidation number under each ion Cross multiply Reduce to simplest form Sodium and chlorine Na Cl Sodium and chlorine Na Cl +1 -1 Sodium and chlorine Na Cl +1 -1 Na 1 Cl 1 Remove any NaCl Reduce “1”s if needed NaCl Iron III and oxygen Fe O Iron III and oxygen Fe O 3+ 2- Iron III and oxygen Fe O 3+ 2- Fe2 O 3 Fe2O3 Cannot be reduced Magnesium Mg O and oxygen Magnesium Mg 2+ O 2- and oxygen Magnesium and oxygen Mg O 2+ 2Mg2 O 2 Mg2O2 Both subscripts can formula is MgO be divided by 2 so the final Silver I and chlorine Magnesium and fluorine Iron II and oxygen Calcium and phosphorus Why did we specify some oxidation numbers, but not others? State the name of the cation. 1. 2. (If using a transition metal, you must state the oxidation number if there is more than one possibility.) State the name of the anion, but change the ending to “ide.” NaCl Sodium chlorine chloride MgO Magnesium oxide K2S Potassium sulfide Fe2O3 Iron is a transition metal so we need to figure out the charge before we can name the compound. We know oxygen is always -2, so there is an overall charge of -6 from the oxygen That means Iron must supply an overall charge of +6 This indicates that iron must have an oxidation number of +3 in this case Fe2O3 Iron III oxide CuS AgCl* H2O (trick!) State the name of the cation. 1. 2. (If using a transition metal, you must state the oxidation number if there is more than one possibility.) Name the anion AgNO3 Silver nitrate CaCO3 Calcium carbonate NH4Cl Ammonium chloride FeSO4 Iron II sulfate 1. 2. 3. 4. Write out each ion. Place oxidation number under each ion Cross multiply Reduce to simplest form Oxyanions are any polyatomic anions that contain oxygen Your book likes to sound fancy! You may not, under any conditions, change the subscripts within the polyatomic ion when balancing the charge. You may only adjust the number of units of each polyatomic ion!! Use parentheses to remind yourself that the units go together and cannot be changed. Potassium K permanganate (MnO4) Potassium K +1 permanganate (MnO4) -1 Potassium K +1 permanganate (MnO4) -1 K 1(MnO4)1 Potassium permanganate is K(MnO4) Calcium Ca hydroxide (OH) Calcium Ca +2 hydroxide (OH) -1 Calcium Ca +2 hydroxide (OH) -1 Ca(OH)2 Ca(OH)2 Ammonium (NH4) phosphate (PO4) Ammonium (NH4) +1 phosphate (PO4) -3 Ammonium (NH4) +1 phosphate (PO4) -3 (NH4)3(PO4) (NH4)3(PO4) Hydrogen peroxide Sodium nitrite Calcium sulfate Aluminum hydroxide A formula unit gives the ration of cations to anions in the ionic compound. A monatomic ion is formed from one atom. Roman numerals indicate the oxidation numbers of any element with more than one oxidation number. Polyatomic ions consist of more than one atom and act as a single unit. To indicate more than one polyatomic ion in a chemical formula, place parentheses around the polyatomic ion and use a subscript outside the parentheses. Relate a formula unit of an ionic compound to its composition Write formulas for ionic compounds and oxyanions. Apply naming conventions to ionic compounds and oxyanions. Metals form crystal lattices and can be modeled as cations surrounded by a “sea” of freely moving valence electrons. Describe a metallic bond Relate the electron sea model the physical properties of metals Define alloys and categorize them into two basic types. Physical property Metal Malleable Electron sea model Delocalized electron Metallic bond Alloy Lattice structures with freely moving electrons Electrons are not firmly attached to any one nucleus, but instead “visit” many nuclei Attraction of a metallic cation for delocalized electrons Freely moving electrons are referred to as “delocalized” (lacking a location) Video - - + + + - - - + + - + - + - - + + - + - - + - - + - + + - + - - + + + + + - - - - - + + + - + + + - + - + - - + - + Melting and Boiling points Vary greatly Most are moderately high melting points and very high boiling points Malleability, ductility and durability Nuclei move relatively free of each other due to the sea of electrons Thermal and electrical conductivity Delocalized electrons quickly move heat from one part of the metal to other parts Delocalized electrons can move in one direction and create a “current.” Hardness and strength The number of delocalized electrons plays a role in the hardness of the metal More delocalized electrons means a harder metal Sometimes d level electrons are delocalized as well as the s resulting in very hard metals. Alloy- a mixture of elements that has metallic properties Characteristics may differ from the “parent” metals Include brass, bronze, 14-carat gold, stainless steel, etc. Substitutional Some of the atoms from one metal are replaced by atoms of the other metal Ex: brass Interstitial Small holes in the lattice are filled by atoms of another element Example: Steel A metallic bond forms when metal cations attract freely moving, delocalized valence electrons. The electron sea model explains the physical properties of metallic solids. Metal alloys are formed when a metal is mixed with one or more other elements. Describe a metallic bond Relate the electron sea model the physical properties of metals Define alloys and categorize them into two basic types.