HIGH SCHOOL SCIENCE

Physical Science 11: Chemical Bonds

WILLMAR PUBLIC SCHOOL

2013-2014 EDITION

C HAPTER 11

Chemical Bonds In this chapter you will: 1. Recognize stable electron configurations. 2.Predict an element’s chemical properties using the number of valence electrons and electron dot diagrams. 3.Describe how ionic bonds form and relate the properties of ionic compounds to the structure of crystal lattices. 4.Describe how covalent bonds form and the attractions that keep atoms together in molecules. 5. Compare polar and nonpolar bonds. 6.Name and determine chemical formulas for ionic and molecular compounds. 7. Describe the structure and strength of bond in metals.

S ECTION 11.1

What are Chemical Bonds?

O BJECTIVES : 1. Recognize stable electron configurations. 2. Predict an element’s chemical properties using the number of valence electrons and electron dot diagrams.

Vocabulary: chemical bond electron dot diagram chemical formula subscript

There is an amazing diversity of matter in the universe, but there are only about 100 elements. How can this relatively small number of pure substances make up all kinds of matter? Elements can combine in many different ways. Did you ever make cupcakes from scratch, like the boy pictured above? You mix together flour, sugar, eggs, and other ingredients to make the batter, put the batter into cupcake papers, and then put them into the oven to bake. The cupcakes that come out of the oven after baking are different from any of the individual ingredients that went into the batter. Like the ingredients that join together to make cupcakes, atoms of different elements can join together to form entirely different substances called compounds. In cupcakes, the eggs and other wet ingredients cause the dry ingredients to stick together. What causes elements to stick together in compounds? Elements form compounds when they combine chemically. Their atoms join together to form molecules, crystals, or other structures. The atoms are held together by chemical bonds. A chemical bond is a force of attraction between atoms or ions. It occurs when atoms share or transfer valence electrons. Valence electrons are the electrons in the outer energy level of an atom. Water (H2O) is an example of a chemical compound. Water molecules always consist of two atoms of hydrogen and one atom of oxygen. Like water, all other chemical compounds consist of a fixed ratio of elements. It doesn’t matter how 2

much or how little of a compound there is. It always has the same composition. The same elements may combine in different ratios. If they do, they form different compounds. Both water (H2O) and hydrogen peroxide (H2O2) consist of hydrogen and oxygen. However, they have different ratios of the two elements. As a result, water and hydrogen peroxide are different compounds with different properties. If you’ve ever used hydrogen peroxide to disinfect a cut, then you know that it is very different from water! Both carbon dioxide (CO2) and carbon monoxide (CO) consist of carbon and oxygen, but in different ratios. There are different types of compounds. They differ in the nature of the bonds that hold their atoms together. The type of bonds in a compound determines many of its properties. Three types of bonds are ionic, covalent, and metallic bonds. An ionic bond is the force of attraction that holds together oppositely charged ions. Ionic bonds form crystals. Table salt contains ionic bonds. A covalent bond is the force of attraction that holds together two nonmetal atoms that share a pair of electrons. One electron is provided by each atom, and the pair of electrons is attracted to the positive nuclei of both atoms. The water molecule represented above contains covalent bonds. A metallic bond is the force of attraction between a positive metal ion and the valence electrons that surround it—both its own valence electrons and those of other ions of the same metal. The ions and electrons form a latticelike structure. Only metals form metallic bonds.

Chemical properties depend on an element’s electron configuration. When the highest occupied energy level of an atom is filled with electrons, the atom is stable and not likely to react. The noble gases have stable electron configurations since their outer shell is full. The chemical properties of an element depend on the number of valence electrons.

An electron dot diagram is a model of an atom in which each dot represents a valence electron. The symbol in the center represents the nucleus and all other electrons in the atom. Electron dot diagrams for carbon and chlorine are shown below. The show paired electrons and unpaired electrons. These valence electrons are available for bonding. The unpaired electrons would like to be in a pair.

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For example, the chemical formula for magnesium chloride is MgCl2. The 2 written to the right and slightly below the symbol of the chlorine is a subscript. Subscripts are used to show the relative numbers of atoms of the element present. If there is only one atom of an element in a formula, no subscript is needed. From the formula, you can tell that there is one magnesium ion for every two chlorine ions in magnesium chlorine. A magnesium atom cannot reach a stable electron configuration by reacting with just one chlorine atom since it has two valence electrons. It must transfer electrons to two chlorine atoms. Elements that do not have complete sets of valence electrons tend to react. By reacting, they achieve electron configurations similar to those of noble gases. You can make a simple salad dressing using just the two ingredients: oil and vinegar. Recipes for oil-and-vinegar salad dressing vary, but they typically include about three parts oil to one part vinegar, or a ratio of 3:1. For example, if you wanted to make a cup of salad dressing, you could mix together 3 cup of oil and 1 cup of vinegar. Chemical compounds also have “ingredients” in a certain ratio. However, unlike oil-and-vinegar salad dressing, a chemical compound always has exactly the same ratio of elements. This ratio can be represented by a chemical formula. A chemical formula is a notation that shows what elements a compound contains and the ratio of the atoms or ions of these elements in the compound.

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Section Review: 1.

Describe when an atom is stable and not likely to react.

2. What does each dot in an electron dot diagram represent? 3. When do elements tend to react? 4. What information is in a chemical formula? 5. What are subscripts used to show? 6. Why must magnesium react with two chlorine atoms to reach a stable electron configuration? 7. The compound sodium sulfide consists of a ratio of one sodium ion (Na+) to two sulfide ions (S-2). Write the chemical formula for this compound. 8. A molecule of sulfur dioxide consists of one sulfur atom (S) and two oxygen atoms (O). What is the chemical formula for this compound? 9. Identify the ratio of atoms in the compound represented by the following chemical formula: N2O5.

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S ECTION 11.2

Some elements achieve stable electron configurations through the transfer of electrons between atoms.

Ionic Bonds O BJECTIVES : 1. Describe how ionic bonds form and relate the properties of ionic compounds to the structure of crystal lattices. 2. Predict an element’s chemical properties using the number of valence electrons and electron dot diagrams.

Look at the electron dot diagram for chlorine. A chlorine atom is missing one valence electron for it to have a full outer shell. Look at the electron dot diagram for sodium. A sodium atom has one valence electron. If sodium were to lose this valence electron, its outer shell would be full and be a stable electron configuration. When sodium reacts with chlorine, an electron is transferred from the sodium atom to the chlorine atom. Each atom ends up with a more stable electron configuration after transferring the electrons than it had before the transfer.

3. Name and determine chemical formulas for ionic and molecular compounds.

Vocabulary: ion

anion

cation

ionic bond

ionization energy

ionic compound

crystals

binary compound

When an atom gains or loses an electron, the number of electrons is no longer equal to the number of protons. The charge is not balance and the charge is not neutral. An ion is an atom that has a positive or negative electric charge. An ion with a negative charge (-) is an anion. An ion with a positive charge (+) is a cation.

polyatomic ions 6

When an anion and cation are close together, a chemical bond forms between the ions. A chemical bond is the force that holds atoms together. An ionic bond is the force that holds cations and anions together. An electron can move to a higher energy level when an atom absorbs energy. The energy allows electrons to overcome the attraction of the positive protons in the nucleus. Cations form when electrons gain enough energy to escape from atoms. It takes energy to remove valence electrons from an atom because the force of attraction between the negative electrons and the positive nucleus must be overcome. The amount of energy needed depends on the element. The amount of energy used to remove an electron is called ionization energy. It varies from element to element. The lower the ionization energy, the easier it is to remove an electron from an atom.

Ionization energies tend to increase from left to right across a period. It takes more energy to remove an electron from a nonmetal than from a metal in the same period. Ionization energies tend to increase from bottom to top in a group. Compounds that contain ionic bonds are ionic compounds. A chemical formula for an ionic compound tells you the ratio of the ions in the compound, but it does not tell you how the ions are arranged in the compound. Solids whose particles are arranged in a lattice structure are called crystals. Lattice structures keep the ions in a fixed position. The repeating pattern of ions in the lattice is like the repeating pattern of designs on wallpaper. Ionic crystals depend on the ratio of ions and their relative size. Crystals are classified into groups based on the shape of their crystals. The properties of an ionic compound can be explained by the strong attractions among ions within a crystal lattice. Ionic compounds have high melting points, are poor conductors of electric current, and shatter when struck with a hammer. 7

The stronger the attraction among the particles, the more kinetic energy the particles must have before they can separate. For an electric current to flow, charge particles must be able to move from one location to another. The ions in a solid crystal lattice have fixed position; however, when solid melts, the lattice breaks apart and the ions are free to flow.

Cu1+ ions to balance an O2ions. CuO is copper (II) oxide because it takes one Cu2+ ions to balance an O2- ions. Polyatomic ions are a covalently bonded group of atoms that has a positive or negative charge and acts as a unit.

When a crystal is struck, negative ions are pushed into positions near negative ions, and positive ions are pushed into position near positive ions. Ions with the same charge repel on another and cause the crystals to shatter. The name of an ionic compound must distinguish the compound from other ionic compounds containing the same elements. The formula of an ionic compound describes the ratio of the ions in the compound. A compound made from only two elements is a binary compound. The naming binary compounds are easy. The name of the cation followed by the name of the anion. The name of the cation is the name of the metal. The name of the anion uses part of the nonmetal with the suffix –ide. For example, if you have sodium and chlorine in a compound, it would be sodium chloride.

If you know the name of the ionic compound you can write its formula. Place the symbol for the cation first, followed by symbol for the anion. Use subscripts to show the ratio of the ions in the compound. Because all compounds are neutral, the total charge on the cations and anions must add up to zero.

Many transition metals form more than one type of ion. When a metal forms more than one ion, the name of the ion contains a Roman numeral to indicate the charge of the ion. For example, Cu2O is copper (I) oxide because it takes two 8

Section Review: 1. How can an atom end up with a more stable electron configuration? 2. What happens when sodium reacts with chlorine? 3. When is it easier to remove an electron from an atom? 4. Across a period, how does the ionization energies increase? 5. In a group, how does the ionization energies increase? 6. Use ionization energy to explain why metals lose electrons more easily than nonmetals. 7. What does the shape of an ionic crystal depend on? 8. How would the properties of an ionic compound be explained? 9. Why do ionic crystals to shatter when struck? 10.What does the name for an ionic compound distinguish? 11.What does the formula for an ionic compound describe? 12.How are binary ionic compounds named? 13.How are transition metals distinguished from one another? 14. How do you write a formula for ionic compound? 9

S ECTION 11.3

Covalent Bonds

In a tennis match, two players keep hitting the ball back and forth. The ball bounces from one player to the other, over and over again. The ball keeps the players moving together on the court. What if the two players represented the nuclei of two atoms and the ball represented valence electrons? What would the back and forth movement of the ball represent? The answer is a covalent bond.

O BJECTIVES : 1. Describe how covalent bonds form and the attractions that keep atoms together in molecules. 2. Compare polar and nonpolar bonds. 3. Name and determine chemical formulas for ionic and molecular compounds.

Vocabulary: covalent bond

molecule

polar covalent bond

nonpolar

A covalent bond is a chemical bond in which two atoms share a pair of valence electrons. Covalent bonds share electrons while ionic bonds transfer electrons. The two atoms that are held together by a covalent bond may be atoms of the same element or different elements. When atoms of different elements form covalent bonds, a new substance, called a covalent compound, results. Water is an example of a covalent compound. A molecule is a neutral group of atoms that are joined together by one or more covalent bonds. The 10

attractions between the shared electrons and the protons in each nucleus hold the atoms together in a covalent bond.

In a molecule of an element, the atoms that form covalent bonds have the same ability to attract an electron. Shared electrons are attracted equally to the nuclei of both atoms. In a molecule of a compound, electrons may not be shared equally. A covalent bond in which electrons are not shared equally is called a polar covalent bond. When atoms form a polar covalent bond, the atom with the greater attraction for electrons has a partial negative charge. The other atom has a partial positive charge.

When two atoms share one pairs of electrons, the bond is called a single bond. When two atoms share two pairs of electrons, the bond is called a double bond. When two atoms share three pairs of electrons, the bond is called a triple bond. In general, elements on the right of the periodic table have a greater attraction for electrons than elements on the left have (except for Noble Gases). In general, elements at the top of a group have greater attraction for electrons than elements at the bottom of a group. Fluorine is on the far right and is at the top of its group, so it has the strongest attraction for electrons and is the most reactive metal. Covalent bonds form because the shared electrons fill each atom’s outer energy level and this is the most stable arrangement of electrons.

Not all polar bonds create a polar molecule. The type of atoms in a molecule and its shape are factors that determine whether a molecule is polar or nonpolar. If the molecule is in a straight line, or linear, with polar bonds, the molecule is non-polar. However, it the molecule is at an angle to each other with polar bonds, the molecule is polar.

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In a molecule compound, there are forces of attraction between molecules. Attractions between polar molecules are stronger than attractions between nonpolar molecules. The covalent bonds of covalent compounds are responsible for many of the properties of the compounds. Because valence electrons are shared in covalent compounds, rather than transferred between atoms as they are in ionic compounds, covalent compounds have very different properties than ionic compounds. • Many covalent compounds, especially those containing carbon and hydrogen, burn easily. In contrast, many ionic compounds do not burn.

compound describe the type and number of atoms in a molecule of the compound. When naming molecular compounds, the general rule is that the most metallic element appears first in the name. If both elements are in the same group, the more metallic element is closer to the bottom of the group. The names of the elements in the compound reflect the actual number of atoms in a molecule. You put a prefix indicating the number of atoms before the element name. If there is only one atom in the first element, the mono is left off. The name of the second element is changed to end in the suffix –ide. For example, N2O4 has two nitrogen atoms and four oxygen atoms. Its name would be dinitrogen tetraoxide.

• Many covalent compounds do not dissolve in water, whereas most ionic compounds dissolve well in water. • Unlike ionic compounds, covalent compounds do not have freely moving electrons, so they cannot conduct 
 electricity. • The individual molecules of covalent compounds are more easily separated than the ions in a crystal, so most 
 covalent compounds have relatively low boiling points. This explains why many of them are liquids or gases at room temperature. Like ionic compounds, molecular compounds have names that identify specific compounds, and formulas that match those names. The name and formula of a molecular

Writing a formula for a molecular compound is easy. Write the symbols for the element in the order the elements appear 12

in the name. The prefixes indicate the number of atoms of each element appear in molecule. The prefixes appears as subscripts in the formulas. If there is no prefix, the number of atoms is one.

Section Review: 1.

How is a covalent bond different from an ionic bond?

2. What keeps the atoms together in a molecule? 3. What do you call it when an atom shared one pair of electrons? 4. Two pairs of electrons? 5. Three pairs of electrons? 6. Which elements have the greatest attraction for electrons? 7. Which atoms become more negative in a polar covalent bond? More positive? 8. How do you determine if a molecule is polar or nonpolar? 9. What does the name and formula for molecular compounds describe? 10. What is the general rule for which element appears first in a molecular compound? 11. How is the second element name changed? 12. What does the prefix hexa- mean? 13. How do you write a formula for molecular compounds? 13

S ECTION 11.4

Metallic Bonds

O BJECTIVE : 1. Describe the structure and strength of bond in metals.

The thick, rigid trunk of the oak tree might crack and break in a strong wind. The slim, flexible trunk of the willow tree might bend without breaking. In one way, metals are like willow trees. They can bend without breaking. That’s because metals form special bonds called metallic bonds. The properties of metal are related to bonds within the metal. There is a way for metal atoms to lose and gain electrons. In a metal, valence electrons are free to move among the atoms, thus becoming a cation with a pool of shared electrons. A metallic bond is the attraction between a metal cation and the shared electrons that surrounds it. Cations in a metal form a lattice that is held in place by strong metallic bonds. Although the electrons are moving among the atoms, the total number of electrons does not change. The valence electrons of metals move freely in this way because metals have relatively low electronegativity, or attraction to electrons. The positive metal ions form a lattice-like structure held together by all the metallic bonds. The more valence electrons an atom can contribute to the shared pool, the stronger the metallic bond.

Vocabulary: metallic bond ductile malleable alloy

The valence electrons surrounding metal ions are constantly moving. This makes metals good conductors of electricity. The lattice-like structure of metal ions is strong but quite flexible. This allows metals to bend without breaking. Metals are both ductile (can be shaped into wires) and malleable (can be shaped into thin sheets).

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Remember that a flow of charge particles is an electric current. A metal has a built-in supply of charged particles that can flow from one location to another, the pool of electrons. The lattice in metals is flexible compared to the rigid lattice of ionic compounds. The metal ions shift their position and shape of the metal changes, but the metal does not shatter. The ions are still held together by the metallic bonds.

other elements in it. The molten solution is then allowed to cool and harden. Alloys generally have more useful properties than pure metals. Alloying one metal with other metal(s) or non metal(s) often enhances its properties. For instance, steel is stronger than iron, its primary element. Unlike pure metals, most alloys do not have a single melting point. Instead, they have a melting range in which the material is a mixture of solid and liquid phases.

Metals such as iron are useful for many purposes because of their unique properties. For example, they can conduct electricity and bend without breaking. However, pure metals may be less useful than mixtures of metals with other elements. For example, adding a little carbon to iron makes it much stronger. This mixture is called steel. Steel is so strong that it can hold up huge bridges, like the one pictured above. Steel is also used to make skyscrapers, cargo ships, cars, and trains. Steel is an example of an alloy. An alloy is a mixture of a metal with one or more other elements. The other elements may be metals, nonmetals, or both. An alloy is formed by melting a metal and dissolving the 15

Section Review: 1.

How do metals achieve a stable electron configuration?

2. What holds metal ions together in a metal lattice? 3. What two important properties of metals can be explained by their structure? 4. What are some useful ways alloys may differ from pure metals?

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