Ionic Covalent Compounds Bonding Chemical bonds are forces

Ionic & Covalent Compounds

Bonding Chemical bonds are forces that cause a group of atoms to behave as a unit. Bonds result from the tendency of a system to seek its lowest possible energy. Bond breaking always requires energy, and bond formation always releases energy.

Types of Bonds The type of bonding depends upon the nature of the atoms that are combined. A metal and a non-metal will form ionic bonds when electrons are transferred from the metal to the non-metal. The resulting attraction between oppositely charged ions forms a stable crystal.

Types of Bonds When metals bond with each other, the valence electrons are shared by the atoms in the entire crystal. The electrons are no longer associated with a specific nucleus, and are free to move throughout the sample.

Lewis Structures, also known as Lewis dot diagrams, show the valence electrons are arranged among the atoms in the molecule. For ionic compounds, it shows the end result when the metal loses its electrons to the non-metal.

Lewis Structures – Ionic Compounds A dot (. ) is used to represent each valence electron. Consider sodium chloride. The dot diagram for the sodium atom is: . : : Na The dot diagram for. chlorine atom is: : the Cl

Lewis Structures – Ionic Compounds : : Note that electrons (dots) representing valence electrons are written singly or in pairs on either side of the atomic symbol, and above and below the symbol. Na. . Cl: It doesn’t matter which side has the unpaired electron.

Lewis Structures – Ionic Compounds Sodium donates its single valence electron to the chorine atom, and chlorine accepts the extra electron. : : . . Na Cl: : : After transfer, both ions have a noble gas electron configuration. Na+ : Cl:

Lattice Energy When metals react with non-metals, a large amount of energy is usually released. These reactions are usually exothermic. The electron transfer from the metal to the non-metal usually requires energy, and is endothermic. The large release of energy comes from forming an ionic crystal.

Lattice Energy The crystal will have oppositely charged ions in contact with each other and like charged ions separated from each other. The alternating charged ions form a crystal lattice. As the lattice forms, large amounts of energy are released.

Lattice Energy

Lattice Energy Lattice energy is defined as the energy required to convert a mole of an ionic solid into its constituent gaseous ions. The greater the lattice energy, the stronger the bonding, and the more stable the compound.

Lattice Energy Based on Coulomb’s Law, the lattice energy increases with smaller ionic size.

Lattice Energy Based on Coulomb’s Law, the lattice energy increases significantly with greater ionic charge. E = constant (q 1)(q 2)/r

Melting Points As ionic crystals melt, the lattice is disrupted, and the ions become free to move, and can conduct electricity.

Ionic Bonding The strength of the attraction between ions increases significantly with increasing ionic charge, and results in high melting points.

Naming Inorganic Compounds 1. Binary Compounds Binary compounds contain only two elements. The elements are either a metal with a non-metal (ionic bonding), or two non-metals (covalent bonding).

Naming Binary Compounds a) Metal + Non-metal: When metals react with non-metals, the metal loses electrons and the non-metal gains electrons. The resulting attraction between oppositely charged ions creates ionic bonds.

Common Ionic Charges The charges of ions of elements in groups 1 A-7 A (the main groups) are usually predictable. Group 1 A metals form +1 ions, group 2 A metals form +2 ions, etc. The non-metals of group 5 A have a -3 charge, those of group 6 A have a -2 charge, and the halogens form ions with a -1 charge.

Typical Ionic Charges

Naming Binary Compounds For example, Na. Cl is called sodium chloride Where “chlor” is the root for the element chlorine.

Naming Binary Compounds Three common transition metals also have only one ionic charge, and are also named the same way. They are: zinc ion (always +2), silver ion (+1) and cadmium ion (+2) Zn. S is zinc sulfide, as “sulf” is the root for sulfur.

Writing Formulas of Binary Compounds have no net charges, so the formulas of ionic compounds must contain equal numbers of positive and negative charges. Magnesium bromide, made from magnesium ion (Mg 2+) and bromide ion (Br 1 -) has the formula Mg. Br 2

Binary Compounds with Variable Charge Metals Most transition metals and the metals on the lower right side of the periodic table can have several ionic charges. The properties of the ion vary greatly with charge, so the charge must be specified in naming the ion or its compounds.

Typical Ionic Charges

Binary Compounds with Variable Charge Metals

Binary Compounds with Variable Charge Metals If an ion has variable charges, you must specify the charge in naming the metal. If an ion has only one charge, it is incorrect to specify its charge.

Naming Fe 2 O 3 n Fe 2 O 3 is an iron oxide, but we must specify the charge of the iron ion. We know each oxide has a -2 charge, so three oxide ions have a total charge of 6. The two iron ions therefore have a charge of +6, with each iron having a charge of +3.

Covalent Bonding Covalent compounds exist as discrete molecules, whereas ionic compounds consist of an aggregate of cations and anions.

Covalent Bonds When two (or more) non-metals form bonds, electrons are shared. The result is a covalent bond. Covalent bonds form because the attraction of electrons for the nuclei in the atoms is greater than the electron-electron repulsion or the nucleus-nucleus repulsion.

Types of Bonds There is usually an optimum bond length or internuclear distance where attractions between electrons and the nuclei are optimized and repulsions are minimized.

Covalent Bonding

Bond Energy Bond formation releases energy, and bond breaking requires energy.

Molecules are neutral combinations of two or more atoms that are bonded together and act as a unit. Molecules may be elements (H 2, O 3, F 2), or compounds containing atoms of different elements bonded together. Molecules typically contain elements that are non-metals.

Molecules Scientists studying the nature of matter focused much of their research in 1800 s on the composition of compounds. Since molecules are much too small to observe, they typically observed the reactions of larger amounts of matter and used mass measurements to develop their theories of matter.

The Law of Definite Proportion n Joseph Proust (1754 -1826) determined the chemical composition of many compounds. He found that a given compound always contains the exact same proportion of elements by mass. This is known as the law of definite proportion. For example, all samples of water contain 88. 8% oxygen by mass, and 11. 2%

The Law of Multiple Proportions This chemical law applies when two (or more) elements can combine to form different compounds. Common examples are carbon monoxide and carbon dioxide, or water and hydrogen peroxide. n John Dalton (1766 -1844) conducted experiments on these types of compounds, and determined that there is a simple relationship between the masses of one element relative to the others. n

The Law of Multiple Proportions n When two elements form a series of compounds, the ratios of the masses of one element that combine with a fixed mass of the other element are always in a ratio of small whole numbers. The meaning of this law is difficult to understand unless it is illustrated using a specific series of compounds.

The Law of Multiple Proportions n Consider the compounds of water and hydrogen peroxide. At this point in history, chemists knew the compounds were different, and that they both contain (or can be broken down into) the elements hydrogen and oxygen. They did not yet know the formulas for either compound, nor was the concept of atoms fully developed.

The Law of Multiple Proportions n Analysis of 100 grams of the compounds produced the following data: Compound water Hydrogen peroxide Mass of Grams of oxygen/100 g hydrogen/10 oxygen/gram of compound 0 g of of hydrogen compound 88. 8 grams O 11. 2 grams H 7. 93 g. O/g. H 94. 06 grams 5. 94 grams H O 15. 8 g. O/g. H

The Law of Multiple Proportions n The Law of Multiple Proportions is illustrated when the numbers in the last column are compared. Compound Grams of 15. 8/7. 93 = 2/1 water Hydrogen peroxide oxygen/gram of hydrogen 7. 93 g. O/g. H 15. 8 g. O/g. H The small whole number ratio suggests that there is twice as much oxygen in hydrogen peroxide as there is in water.

The Law of Multiple Proportions n The key feature is that small whole numbers are generated. The results support the hypothesis that molecules consist of various combinations of atoms, and that atoms are the smallest unit of matter. The ratio doesn’t produce fractions, since there is no such thing as a fraction of an atom. n For the example cited, we would propose that hydrogen peroxide contains twice as many oxygen atoms/hydrogen atoms than does water. We cannot, however, determine the actual formula of either compound.

Empirical Formulas The studies of water and hydrogen peroxide lead to empirical formulas. These are based on experiment, and represent to simplest way of expressing the ratio of atoms in a compound. The early scientists analyzed new chemical compounds to determine their composition and chemical formulas. Modern analytical laboratories still provide this service.

Molecular Formulas Molecular formulas show the exact number of each type of atom in a molecule. For example, hydrogen peroxide has a molecular formula of H 2 O 2. Its empirical formula shows that there is one hydrogen for every oxygen, so it is OH or HO. Neither of these formulas provides the structure or arrangement of the bonds in the molecule.

Structural Formulas Structural formulas provide the arrangement of atoms in the molecule. The structural formula for hydrogen peroxide is: H-O-O-H This formula shows the arrangement of the atoms, but doesn’t show bond angles or the shape of the molecule.

Naming Covalent Binary Compounds When two non-metals form a compound, they share electrons, rather than transfer them. The resulting bond is called a covalent bond. The naming of these compounds is fairly simple. The first element is named first, and the second element is named as the root + ide. Prefixes are used to indicate the number of each atom present.

Naming Covalent Binary Compounds These prefixes are used only for compounds containing two nonmetals. The prefix mono is never used for the first element in the compound.

Naming Covalent Binary Compounds n The prefix mono is never used for the first element. CO 2 is carbon dioxide. n If the prefix ends in an a or o, and the element that follows begins with a vowel, the last letter of the prefix is usually dropped. N 2 O 5 is called dintrogen pentoxide (and not pentaoxide).

Naming Covalent Binary Compounds Note that these prefixes are only used for binary covalent compounds. It is incorrect to use them for compounds containing a metal and a non-metal.

Naming Covalent Binary Compounds There are some compounds of metalloids or metals in very high oxidation states that are sometimes named using this system.

Naming Covalent Binary Compounds For example, compounds such as Sn. Cl 4 or Pb. Cl 4 are covalent in nature, and not ionic solids. They may be called tin(IV) chloride or tin etrachloride [or lead(IV) chloride or lead tetrachloride.

Binary Compounds with Hydrogen With metals, hydrogen can form ionic compounds in which the hydrogen has a 1 ionic charge. These compounds are named like any binary ionic compound. Na. H is sodium hydride Ca. H 2 is calcium hydride

Binary Compounds with Hydrogen With non-metals, the bonding is covalent. Hydrogen never forms a positive ion in nature. Many of the compounds containing hydrogen have common names that do not follow the usual nomenclature rules.

Binary Compounds with Hydrogen Examples include: water H 2 O ammonia NH 3 phosphine PH 3 hydrogen sulfide H 2 S Note that the order in which the elements are written is also irregular.

Binary Compounds with Hydrogen- Acids Hydrogen also forms binary compounds that act as acids in water. These compounds dissociate in water to donate a proton to water. HCl(g) + H 2 O(l) H 3 O+(aq) + Cl– (aq) hydrogen chloride hydronium

Naming Binary Acids The naming of the pure compound and its aqueous acid solution differ. HCl is a gas called hydrogen chloride. HCl(aq) is an acid called hydrochoric acid.

Naming Binary Acids n Name the following acids: H 2 S(aq) , HBr(aq)

Organic Nomenclature Compounds containing carbon, hydrogen and sometimes oxygen, nitrogen, sulfur and the halogens, have a separate system of nomenclature.

Unusual Ions Mercury forms two ions, mercury(I) and mercury(II). The mercury(I) ion is polyatomic, and exists as two mercury(I) ions bonded together. Its formula is Hg 22+. Oxygen in compounds usually exists as the oxide ion, O 2 -. Oxygen also exists as the peroxide ion, O 22 -, with each oxygen having a -1 charge.

Naming Polyatomic Ions There are many ions, such as sulfate or nitrate, that contain more than one element. Many of these ions contain oxygen and a non-metal. These ions can be found in a group of acids called the oxy acids (such as sulfuric acid, nitric acid, etc. ).

Polyatomic Ions The bonding within these polyatomic ions (such as nitrate, sulfate and phosphate) is covalent. The ionic charge results from the loss of one or more H ions to water, resulting in a negative charge on the anion formed. In water, the covalently bonded hydrogen is donated to water, forming hydronium ions and the corresponding anion.

Naming the Oxy Acids The easiest way to learn the names of the ions is to memorize a short list of oxy acid names and their formulas. The names of the ions are derived from the names of the acids. Keep in mind that the acids must be aqueous solutions.

Common Oxy Acids Acid HNO 3 Name Nitric acid H 2 SO 4 Sulfuric acid HCl. O 3 acid) Chloric acid (or iodic or bromic H 3 PO 4 Phosphoric acid H 2 CO 3 Carbonic acid

Naming Complex Ions Once the list of acids is learned, the names of other acids and ions can be derived. Removal of the hydrogens in the acid as H+ ions results in ions that end in ate. HNO 3 minus one H+ ion gives NO 31 -, the nitrate ion. The oxy acids that end in ic, produce ions that end ate.

Naming Complex Ions Sulfuric acid is H 2 SO 4. Removing two H+ ions produces SO 42 -, the sulfate ion. Keep in mind that the formula of the ions must include the charge. If only one of the H+ ions is removed from sulfuric acid, HSO 41 - is produced. This is called the hydrogen sulfate ion, also commonly known as the bisulfate ion.

Naming Complex Ions Carbonic acid, H 2 CO 3, produces two ions: HCO 31 -, the hydrogen carbonate or bicarbonate ion and CO 32 -, the carbonate ion

Naming Complex Ions Some of the oxy acids previously listed also exist with one more oxygen in the formula. HCl. O 3, HBr. O 3 and HIO 3 , in aqueous solution are chloric, bromic and iodic acid respectively. Adding an oxygen to the formulas provides the formulas for the per root ic acid. HCl. O 4 is perchloric acid. The ion, Cl. O 41 - is the perchlorate ion.

Naming Complex Ions Several of the oxy acids listed previously can have one less oxygen atom in the formula. These acids have names that end in ous, and ions that end in ite. HNO 3 is nitric acid. HNO 2(aq) is nitrous acid. The ion NO 21 - is the nitrite ion.

Naming Complex Ions Sulfuric acid, phosphoric acid, chloric, bromic and iodic acids all can have one less oxygen atom. The acids are sulfurous acid, phosphorous acid, chlorous acid, bromous acid and iodous acid. The ions are called sulfite, phosphite, chlorite, bromite and iodite ion.

Naming Complex Ions The halogen oxy acids HCl. O 3, HBr. O 3, and HIO 3 also exist with two less oxygen atoms in the formula. The name of the resulting acid has the name hypo root ous acid. HCl. O(aq) is hypochlorous acid, and Cl. O 1 - is the hypochlorite ion.

Naming Complex Ions n If you memorize the list of acids ending in ic, you can derive the names and formulas for many other acids and ions. Acid HNO 3 H 2 SO 4 HCl. O 3 acid) H 3 PO 4 H 2 CO 3 Name Nitric acid Sulfuric acid Chloric acid (or iodic or bromic Phosphoric acid Carbonic acid

Naming Complex Ions In naming the ions from the acids on the list, remember that ic ate. n If there is one additional oxygen atom, the acid has the name per root ic, and the ion has the name per root ate. n If there is one less oxygen atom, the acid has a name ending in ous. The ions will have names ending in ite. (ous ite) n

Naming Complex Ions n If an acid has two less oxygen atoms than the “ic” list, its name has the form hypo root ous. The ion will have the name hypo root ite.

Other Common Formulas CH 3 COOH CH 3 COO 1 NH 3 NH 4+ OH 1 H 3 O + Mn. O 41 Cr. O 42 Cr 2 O 72 - Acetic acid Acetate ion Ammonia Ammonium ion Hydroxide ion Hydronium ion Permanganate ion Chromate ion Dichromate ion

Percent Composition A chemical formula can be used to calculate the percent composition of a compound. Likewise, the percent composition can be used to determine the empirical formula of a compound. This is extremely useful when trying to determine the formula of a new, or unknown compound.

Chemical Composition Usually, the compound is combusted in the presence of oxygen. Any carbon in the compound is collected as carbon dioxide (CO 2), and any hydrogen is collected as water (H 2 O).

Chemical Composition Similar techniques exist to analyze for other elements. The formula obtained for the compound is the simplest whole number ratio of the elements in the compound, or the empirical formula. It may differ from the actual formula. For example, hydrogen peroxide is H 2 O 2, but chemical analysis will provide an empirical formula of HO.

Percent Composition To calculate the % composition of a known compound, you determine the total mass of the molecule, and the mass due to each of the elements in the compound. % by mass of element A = total mass of A (100%) molecular mass

% Composition Problem: Calculate the percent composition of ammonia.

Determining Formulas It is generally more useful to obtain percent composition data (usually from a laboratory), and determine the empirical formula of a compound. This will be the simplest whole number ratio of the elements, and provides no information about the structure of the compound.

Determining Empirical Formulas If given % composition: 1. Assume a quantity of 100 grams of the compound. 2. Determine the number of moles of each element in the compound by dividing the grams of each element by the appropriate atomic mass. 3. To simplify the formula into small whole numbers, divide the moles of each element by the smallest number of moles.

Determining Empirical Formulas 4. If necessary, multiply each number of moles by a factor that produces whole number subscripts. 5. If you know the approximate molar mass of the compound, determine the molecular formula.

% Composition Problem n Caffeine contains 49. 48% carbon, 5. 15% hydrogen, 28. 87% nitrogen, and 16. 49% oxygen. The compound has a molar mass of 194. 2. Determine the empirical and molecular formula of caffeine.