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Chapter 3

Chapter 3. Compounds: Putting Particles Together. Chapter Outline. 3.1 The Octet Rule 3.2 In Search of an Octet Part 1: Ion Formation 3.3 Ionic Compounds—Electron Give and Take 3.4 In Search of an Octet Part 2: Covalent Bonding 3.5 Getting Covalent Compounds into Shape

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Chapter 3

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  1. Chapter 3 Compounds:Putting Particles Together

  2. Chapter Outline 3.1 The Octet Rule 3.2 In Search of an Octet Part 1: Ion Formation 3.3 Ionic Compounds—Electron Give and Take 3.4 In Search of an Octet Part 2: Covalent Bonding 3.5 Getting Covalent Compounds into Shape 3.6 Electronegativity and Molecular Polarity Chapter 3

  3. 3.1 The Octet Rule • What drive’s most elements to form compounds? • Consider Group 8A elements—the noble gases. • Noble gases are chemically unreactive and highly stable. • If an element is more stable, it is less likely to react with another element, and if an element is less stable, it is likely to react. Chapter 3

  4. 3.1 The Octet Rule, Continued • The stability of an element and its reactivity is inversely related. • What makes Group 8A elements more stable and less reactive than other elements? • Recall that the group number indicates the number of valence electrons for any element in that group. Each noble gas has eight valence electrons (except helium, which has two). Chapter 3

  5. 3.1 The Octet Rule, Continued • An electron arrangement of eight valence electrons is unusually stable. Chemists do not completely understand why this arrangement is so stable. • The octet rule: Most atoms react to achieve a total of eight valence electrons in their valence shell. Chapter 3

  6. 3.2 In Search of an Octet Part 1: Ion Formation • Atoms gain or lose electrons to achieve a valence octet. • An atom contains an even number of electrons and protons, which means it is electrically neutral. • If an atom gains or loses electrons to form a valence octet, the new particle would have an unequal number of protons and electrons and would have a net charge. Chapter 3

  7. 3.2 In Search of an Octet Part 1: Ion Formation, Continued • An ion is an atom that has gained or lost electrons and has a charge. • Group 7A elements, like chlorine, have seven valence electrons, gain an electron to have an octet, and form an ion with a 1- charge (ions have 17 protons and 18 electrons). • NOTE: When chlorine forms an ion it has the same number of valence electrons and total electrons as argon. Therefore, the chlorine ion is said to be isoelectronic(iso, a prefix meaning “same”) to argon. Chapter 3

  8. 3.2 In Search of an Octet Part 1: Ion Formation, Continued An anion is any atom that contains a net negative charge given that it has an unequal number of protons and electrons. Chapter 3

  9. 3.2 In Search of an Octet Part 1: Ion Formation, Continued • Group 1A elements, like sodium, have one valence electron. In order to have an octet, the elements in this group must lose one electron to form ions with a 1+ charge (ions have 11 protons and 10 electrons). • NOTE: When sodium forms an ion it has the same number of valence electrons and total electrons as argon. Therefore, the sodium ion is isoelectronic to neon. Chapter 3

  10. 3.2 In Search of an Octet Part 1: Ion Formation, Continued A cation is any atom that contains a net positive charge, given that it has an unequal number of protons and electrons. Chapter 3

  11. 3.2 In Search of an Octet Part 1: Ion Formation, Continued Trends in Ion Formation • How do we determine which atoms will gain electrons to form anions and which atoms will lose electrons to form cations? • Metals lose electrons to form cations and nonmetals gain electrons to form anions. • The main group metals in Groups 1A, 2A, and 3A form cations. Main group nonmetals in Groups 5A, 6A, and 7A form anions. Chapter 3

  12. 3.2 In Search of an Octet Part 1: Ion Formation, Continued The main group metals form cations with charges of 1+ to 3+, whereas the main group nonmetals form anions with charges of 1- to 3-. Chapter 3

  13. 3.2 In Search of an Octet Part 1: Ion Formation, Continued • Group 1A elements lose one electron to form cations with a 1+ charge. Group 2A elements lose two electrons to form cations with a 2+ charge. Group 3A elements give up three electrons to form cations with a 3+ charge. • With main group metals, the group number indicates the charge of the cation that is formed. Chapter 3

  14. 3.2 In Search of an Octet Part 1: Ion Formation, Continued • The charge on the main group elements can be determined by subtracting eight from the group number. For example, for Group 6A elements, the ions formed have a 2- charge (6 – 8 = 2-). • The rules for predicting charges on ions formed by the main group elements are as follows: Chapter 3

  15. 3.2 In Search of an Octet Part 1: Ion Formation, Continued • Predicting the charge of metals (other than the main group metals) is not as simple. • Transition metals form more than one ion. For example, iron will form the Fe2+ and Fe3+ cation. Copper will form the Cu+ and Cu2+ cation. • Another group of common ions whose charge cannot be determined by the periodic table are those formed when two or more elements combine to form an ion. Chapter 3

  16. 3.2 In Search of an Octet Part 1: Ion Formation, Continued Polyatomic ionsare a special group of ions formed when two or more nonmetal elements interact with one another. The ions formed have a single charge. Chapter 3

  17. 3.2 In Search of an Octet Part 1: Ion Formation, Continued Naming Ions • To distinguish between an element and its ion, a different name is given to the ion. • To name metal ions, add the word ion to the name of the metal, so Na+ becomes the sodium ion. • Transition metals form more than one ion. A roman numeral in parentheses following the name of the metal indicates the charge. F2+ is iron(II) and Fe3+ is iron(III). Chapter 3

  18. 3.2 In Search of an Octet Part 1: Ion Formation, Continued • To name nonmetal ions, the suffix –ide replaces the last few letters of the name of the element. The anion formed from chlorine becomes the chloride ion. The anion formed from oxygen is the oxide ion. • To name polyatomic ions, keep in mind thatmost polyatomic ions end in -ate, but the ending –ite is used for names of related ions that have one less oxygen atom. The hydroxide (OH-), hydronium (H3O+), and cyanide (CN-) ions are exceptions to this rule. Chapter 3

  19. 3.2 In Search of an Octet Part 1: Ion Formation, Continued Important Ions in the Body • Ions in the body perform important functions. • Main cations in the body are Na+, K+, Ca2+,and Mg2+. They are important in maintaining solution concentrations inside and outside cells. • Main anions in the body are Cl-, HCO3-, and HPO42-. These help maintain charge neutrality between blood and fluids in the cell. Chapter 3

  20. 3.2 In Search of an Octet Part 1: Ion Formation, Continued Chapter 3

  21. 3.3 Ionic Compounds—Electron Give and Take Cations do not form unless anions form and vice versa. When metals and nonmetals combine, electrons are transferred between the atoms. Ion formation is shown as follows: Chapter 3

  22. 3.3 Ionic Compounds—Electron Give and Take, Continued • Opposite charges attract, so any cation and anion formed are strongly attracted to one another. • An ionic bond isformed when a cation is attracted to an anion. The resulting combination of cation and anion is called an ionic compound. Chapter 3

  23. 3.3 Ionic Compounds—Electron Give and Take, Continued Formation of table salt, sodium chloride, is shown as follows: Chapter 3

  24. 3.3 Ionic Compounds—Electron Give and Take, Continued Formulas of Ionic Compounds When ions are attracted to one another to form ionic compounds, the overall charge is zero. Chapter 3

  25. 3.3 Ionic Compounds—Electron Give and Take, Continued • Na+ and Cl- interact to form a neutral ionic compound, NaCl. The formula is written with no charges shown. • When a formula is written, the number of each ion in the compound is represented as a subscript. • Just as a charge of 1+ on an ion is represented as +, the “1” is understood in the formula. Sodium chloride is written as NaCl, not as Na1Cl1. Chapter 3

  26. 3.3 Ionic Compounds—Electron Give and Take, Continued How do you write the formula if one ion has a charge greater than the other? • Cations and anions combine so the total charge on the ionic compound has no net charge. • Consider calcium ions with a charge of 2+ and fluorine ions with a charge of 1-. • A calcium atom reacts with two fluorine atoms giving up two electrons, one to each of the fluorine atoms. • One Ca2+ ion is formed and two F- ions are formed. Chapter 3

  27. 3.3 Ionic Compounds—Electron Give and Take, Continued The overall charge resulting from the ionic compound formed when one calcium atom reacts with two fluorine atoms is shown as: Chapter 3

  28. 3.3 Ionic Compounds—Electron Give and Take, Continued How do we determine the formula for the ionic compound formed with Al3+ reacts with O2-? • Use the method known as the crossover approach. • Using the crossover approach, the charge number of the cation becomes the subscript of the anion, and the charge number of the anion becomes the subscript of the cation. Chapter 3

  29. 3.3 Ionic Compounds—Electron Give and Take, Continued Naming Ionic Compounds When naming ionic compounds, combine the names of the ions involved after dropping the word ion from the name of the individual ions. Chapter 3

  30. 3.4 In Search of an Octet Part 2:Covalent Bonding • To form an octet in the valence shell, some atoms transfer or accept electrons to form ions, and some atoms chemically combine with other atoms to share valence electrons. • A covalent bond isformed by the sharing of electrons between two nonmetals. • A covalent compound isformed when nonmetals share electrons.The smallest unit of a covalent compound is called a molecule. Chapter 3

  31. 3.4 In Search of an Octet Part 2:Covalent Bonding, Continued Covalent Bond Formation • The rule of thumb in covalent bond formation is that the number of covalent bonds that an atom will form equals the number of electrons that it needs to complete its octet. • The electron dot symbol (elemental symbol plus a dot for each valence electron) for an atom is used to determine covalent bond formation. Chapter 3

  32. 3.4 In Search of an Octet Part 2:Covalent Bonding, Continued Below is the list of steps for the construction of an electron dot symbol: • Write the symbol of the element. • Place one valence electron on the top, bottom, right, and left of the element symbol. • Continue placing the remaining valence electrons around the element symbol by forming ion pairs with the first four electrons. • Chlorine has seven valence electrons and its electron dot symbol is: Chapter 3

  33. 3.4 In Search of an Octet Part 2:Covalent Bonding, Continued • One chlorine atom can form a covalent bond with another chlorine atom. • Chlorine needs only one electron to complete its octet, so it shares an electron from another chlorine atom. • A bonding pair is a shared pair of electrons. The remaining valence electrons are known as the lone pair of electrons. Chapter 3

  34. 3.4 In Search of an Octet Part 2:Covalent Bonding, Continued • The bonding and lone pair of electrons are shown for the newly formed molecule Cl2. • Below are the electron dot symbols of the main elements found in living systems: Chapter 3

  35. 3.4 In Search of an Octet Part 2:Covalent Bonding, Continued • Carbon, the element of life on Earth, has four unpaired valence electrons. • Carbon can form: • Four single bonds (a sharing of two electrons) • One double and two single bonds, or two double bonds (a sharing of four electrons) • One single bond and one triple bond (a sharing of six electrons) Chapter 3

  36. 3.4 In Search of an Octet Part 2:Covalent Bonding, Continued Chapter 3

  37. 3.4 In Search of an Octet Part 2:Covalent Bonding, Continued Formulas and Structures of Covalent Compounds • A molecular formula identifies all the components in a molecule of a covalent compound. • Table sugar, sucrose (molecular formula C12H22O11) has 12 carbon atoms, 22 hydrogen atoms, and 11 oxygen atoms. • Molecular formula only tells us the number of atoms in a molecule not how they are joined or what the structure is. Chapter 3

  38. 3.4 In Search of an Octet Part 2:Covalent Bonding, Continued • Consider methane, which has a molecular formula of CH4. It contains one carbon atom and four hydrogen atoms. • To determine how atoms are connected, start by drawing the electron dot symbols for all the atoms involved. Chapter 3

  39. 3.4 In Search of an Octet Part 2:Covalent Bonding, Continued • Carbon has four unpaired valence electrons and needs four to complete its octet. • Each hydrogen atom has one valence electron and needs two electrons for completion (an exception to the octet rule), therefore hydrogen can form one single bond. • Carbon can form one single bond with each of the hydrogen atoms in its molecular formula. This representation is called a Lewis structure. Chapter 3

  40. 3.4 In Search of an Octet Part 2:Covalent Bonding, Continued Naming Covalent Compounds • A binary compound is acovalent compound composed of only two elements. • The following are rules for naming binary compounds: Step 1. Name the first element in the formula. Step 2. Name the second element and change the ending to -ide. Step 3. Designate the number of each element present using a Greek prefix. Chapter 3

  41. 3.4 In Search of an Octet Part 2:Covalent Bonding, Continued Indicating the number of each element present is important. For example, carbon and oxygen can form both CO (carbon monoxide) and CO2 (carbon dioxide). The prefix mono- for carbon is understood. Chapter 3

  42. 3.5 Getting Covalent Compounds into Shape • To understand molecules it is important to identify the three-dimensional shapes of each molecule. • The Lewis structure of methane appears flat on paper, which implies the four hydrogen bond angles are 90o. • Electrons are negatively charged and repel each other, so the clouds of negative charge want to be as far away from each other as possible. Chapter 3

  43. 3.5 Getting Covalent Compounds into Shape, Continued When the electrons obtain maximum distance from each other, the three-dimensional structure of methane shows a molecule whose atoms achieve a tetrahedral shape, with bond angles of 109.5o between the clouds of electrons. Chapter 3

  44. 3.5 Getting Covalent Compounds into Shape, Continued • The valence-shell electron-pair repulsion (VSEPR) model is used to predict the shape of a molecule based on the number of electron-charged clouds on a given atom. • VSEPR states that in the valence shell of an atom, the charge clouds formed by groups of electrons will arrange themselves to be as far away from each other as possible in order to reduce repulsions. Chapter 3

  45. 3.5 Getting Covalent Compounds into Shape, Continued Determining the Shape of a Molecule The following are steps used to determine the shape of a molecule using VSEPR: Step 1. Determine the central atom of a molecule. The central atom is the one the shape will be centered around. It is represented as A in the VSEPR form. Step 2. Determine the number of charge clouds around the central atom. Single and multiple bonds are counted as one charge cloud. Charge clouds in bonds are called bonding clouds. They are represented by B in the VSEPR form. Chapter 3

  46. 3.5 Getting Covalent Compounds into Shape, Continued • Lone pair of electrons are counted as one charge cloud and affect the shape of the molecule. • Nonbonding clouds are charge clouds from lone pairs of electrons and are represented as N in the VSEPR form. Chapter 3

  47. 3.5 Getting Covalent Compounds into Shape, Continued One example is formaldehyde, which has a VSEPR form of AB3—a central atom with three bonding charge clouds. Its three-dimensional structure is shown as: Chapter 3

  48. 3.5 Getting Covalent Compounds into Shape, Continued Another example is ammonia, which has a VSEPR form of AB3N—a central atom with three bonding charge clouds and one nonbonding charge cloud. Its three-dimensional structure is shown as: Chapter 3

  49. 3.5 Getting Covalent Compounds into Shape, Continued The following lists molecular shapes associated with the VSEPR forms. Chapter 3

  50. 3.5 Getting Covalent Compounds into Shape, Continued Nonbonding Pairs and Their Effect on Molecular Shape • Nonbonding pairs affect the shape of a molecule. • A molecule like methane has a central atom with four atoms around it and assumes the tetrahedral shape. • Ammonia is a molecule with three atoms and a nonbonding pair around central atoms. Its shape changes to pyramidal. Chapter 3

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