CHEMISTRY The Central Science 9th Edition

1. CHEMISTRYThe Central Science 9th Edition Chapter 2Atoms, Molecules, and Ions

2. Dalton’s law of multiple proportions: When two elements form different compounds, the mass ratio of the elements in one compound is related to the mass ratio in the other by a small whole number. Atomic theory: Each element is composed of tiny particles called atoms All atoms of a given element are identical. In chemical reactions, the atoms are not changed. Compounds are formed when atoms of more than one element combine. The Atomic Theory of Matter

3. The Discovery of Atomic Structure • Atoms are the building blocks of matter. • The ancient Greeks were the first to postulate that matter consists of indivisible constituents. • Later scientists realized that the atom consisted of charged entities.

4. The Modern View of Atomic Structure • The atom consists of positive, negative, and neutral entities (protons, electrons, and neutrons). • Protons and neutrons are located in the nucleus of the atom, which is small. Most of the mass of the atom is due to the nucleus. • There can be a variable number of neutrons for the same number of protons. Isotopes have the same number of protons but different numbers of neutrons. • Electrons are located outside of the nucleus. Most of the volume of the atom is due to electrons.

5. The Atom

6. Class Practice Problem • The diameter of a U.S. penny is 19mm. The diameter of a copper atom, by comparison, is only 2.6 angstroms (Å). How many copper atoms could be arranged side by side in a straight line across the diameter of a penny?

7. Atomic Weights • The Atomic Mass Scale • 1H weighs 1.6735 x 10-24 g and 16O 2.6560 x 10-23 g. • We define: mass of 12C = exactly 12 amu. • Using atomic mass units: • 1 amu = 1.66054 x 10-24 g • 1 g = 6.02214 x 1023 amu

8. Atomic Number, Mass Number, and Isotopes • Atomic number (Z) = number of protons in the nucleus. • Mass number (A) = total number of nucleons in the nucleus (i.e., protons and neutrons). • By convention, for element X, we write ZAX. • Isotopes have the same Z but different A. • We find Z on the periodic table.

9. Class Practice Problem • How many protons, neutrons, and electrons are in an atom of 197Au? • Hydrogen has three isotopes, with mass numbers 1, 2, and 3. Write the complete chemical symbol for each of them.

10. Atomic Weights • Average Atomic Masses • Relative atomic mass: average masses of isotopes: • Naturally occurring C: 98.892 % 12C + 1.108 % 13C. • Average mass of C: • (0.98892)(12 amu) + (0.0108)(13.00335) = 12.011 amu. • Atomic weight (AW) is also known as average atomic mass (atomic weight). • Atomic weights are listed on the periodic table.

11. Arrangement of the Periodic Table • The Periodic Table is used to organize the 114 elements in a meaningful way. • As a consequence of this organization, there are periodic properties associated with the periodic table.

12. The Periodic Table

13. Reading the Periodic Table • Columns in the periodic table are called groups (numbered from 1A to 8A or 1 to 18). • Rows in the periodic table are called periods. • Metals are located on the left hand side of the periodic table (most of the elements are metals). • Non-metals are located in the top right hand side of the periodic table. • Elements with properties similar to both metals and non-metals are called metalloids and are located at the interface between the metals and non-metals.

14. Properties of the Periodic Table • Some of the groups in the periodic table are given special names. • These names indicate the similarities between group members: • Group 1A: Alkali metals. • Group 2A: Alkaline earth metals. • Group 6A: Chalcogens. • Group 7A: Halogens. • Group 8A: Noble gases.

15. Molecules and Molecular Compounds • Molecules are assemblies of two or more atoms bonded together. • Each molecule has a chemical formula. • The chemical formula indicates • which atoms are found in the molecule, and • in what proportion they are found. • Compounds formed from molecules are molecular compounds. • Molecules that contain two atoms of the same element bonded together are called diatomicmolecules.

16. Molecules and Molecular Compounds Example of Diatomic Molecules

17. Molecules and Molecular Compounds • Molecular and Empirical Formulas • Molecular formulas • give the actual numbers and types of atoms in a molecule. • Examples: H2O, CO2, CO, CH4, H2O2, O2, O3, and C2H4.

18. Molecules and Molecular Compounds • Most molecular substances that we will study in this class contain only nonmetals.

19. Molecules and Molecular Compounds • Molecular and Empirical Formulas • Empirical formulas • give the relative numbers and types of atoms in a molecule. • That is, they give the lowest whole number ratio of atoms in a molecule. • Examples: H2O, CO2, CO, CH4, HO, CH2.

20. Molecules and Molecular Compounds • Molecular and empirical formulas do not show how atoms are arranged when bonded together.

21. Molecules and Molecular Compounds • Picturing Molecules • Molecules occupy three dimensional space. • However, we often represent them in two dimensions. • The structural formula gives the connectivity between individual atoms in the molecule. • The structural formula may or may not be used to show the three dimensional shape of the molecule. • If the structural formula does show the shape of the molecule, then either a perspective drawing, ball-and-stick model, or space-filling model is used.

22. Molecules and Molecular Compounds Representing Structure in Molecules Accurately represents the angles at which molecules are attached.

23. Class Practice Exercise The structural formula of propaneand butane is What is the chemical and empirical formula for these molecules? H H H H H H H C H C C H C H C C C H H H H H H H H

24. Ions and Ionic Compounds • When an atom or molecule loses electrons, it becomes positively charged. • For example, when Na loses an electron it becomes Na+. • Positively charged ions are called cations. • When an atom or molecule gains electrons, it becomes negatively charged. • For example when Cl gains an electron it becomes Cl-. • Negatively charged ions are called anions. • An atom or molecule can lose more than one electron. • When molecules loose electrons, polyatomic ions are formed.

25. Ions and Ionic Compounds • In general: metal atoms tend to lose electrons to become cations; nonmetal ions tend to gain electrons to form anions. • Predicting Ionic Charge • The number of electrons an atom loses is related to its position on the periodic table.

26. Ions and Ionic Compounds Predicting Ionic Charge

27. Ions and Ionic Compounds • Element Bonding • The majority of chemistry involves the transfer of electrons between species. • Example: • To form NaCl, the neutral sodium atom, Na, must lose an electron to become a cation: Na+. • The electron cannot be lost entirely, so it is transferred to a chlorine atom, Cl, which then becomes an anion: Cl-. • The Na+ and Cl- ions are attracted to form an ionic NaCl lattice which crystallizes. • NaCl is an example of an Ionic compound (consisting of positive and negatively charged atoms)

28. Ions and Ionic Compounds Crystal Structure of NaCl

29. Ions and Ionic Compounds • Ionic Compounds • Important: note that there are no easily identified NaCl molecules in the ionic lattice. Therefore, we cannot use molecular formulas to describe ionic substances. • Writing the empirical formulas for ionic compounds: • you need to know the ions of which it is composed. • The formula must reflect the electrical neutrality of the compound • the total positive charge must equal the total negative charge • Example: Consider the formation of Mg3N2: • Mg loses two electrons to become Mg2+; • Nitrogen gains three electrons to become N3-. • For a neutral species, the number of electrons lost and gained must be equal.

30. Ions and Ionic Compounds • Writing the Empirical Formula • However, Mg can only lose electrons in twos and N can only accept electrons in threes. • Therefore, Mg needs to lose 6 electrons (2  3) and N gain those 6 electrons (3  2). • I.e., 3Mg atoms need to form 3Mg2+ ions (total 3  2+ charges) and 2 N atoms need to form 2N3- ions (total 2  3- charges). • Therefore, the formula is Mg3N2.

31. Controversy in Naming Inorganic/Organic Compounds Organic compounds contain carbon. Inorganic compounds don't. This definition is often given but is no help at all. What do we make of carbon dioxide, sodium cyanide, baking soda (sodium bicarbonate), ...? Organic compounds contain carbon-hydrogen bonds. Inorganic compounds don't. This is a much better definition, allowing us to call sodium acetylide "organic" but calcium carbide "inorganic," but it doesn't always work. Inorganic compounds contain metal atoms. Organic compounds don't. This doesn't really work any too well either. Even leaving the huge field of organometallic chemistry out of the running, are we really going to call soap (sodium salts of fatty acids) or the lipid bilayers forming cell membranes (again, salts of long-chain organic acids) "inorganic"??? An organic compound is whatever an organic chemist says it is; an inorganic compound is whatever an inorganic chemist says it is. Reproduced from http://www.madsci.org/posts/archives/dec2000/975719013.Ch.r.html

32. Naming Inorganic Compounds • Naming of compounds, nomenclature, is divided into organic compounds (those containing C, usually in combination with hydrogen) and inorganic compounds (the rest of the periodic table). • Naming Ionic Compounds • Based on the names of the ions of which they are composed. • Example, NaCl is called sodium chloride (based on Na+ and Cl- ions). • The cation is written first and the anion is written last. • Ions may be monoatomic or polyatomic. • Vast majority of monoatomic cations are made from metals. • These ions take the name of the element itself.

33. Naming Inorganic Cations • Cations formed from a metal have the same name as the metal. • Example: Na+ = sodium ion. • If the metal can form more than one cation, then the charge is indicated by a Roman numeral in parentheses in the name. • Examples: Cu+ = copper(I); Cu2+ = copper(II) (Page 61). • Most of the elements that can form more than one cation are the • transition metals (3B to 2B). • Or placing ous or ic at the end of the name to indicate the lower • and higher, respectively, charged cation. • Cations formed from non-metals (end in -ium). • Example: NH4+ ammonium ion.

34. Some Common Cations

35. Naming Inorganic Anions • Monoatomic anions (with only one atom) are named by dropping the ending of the name and replacing with -ide. • Example: Cl- is chloride ion. • Polyatomic anions (with many atoms) containing oxygen end in -ate or -ite. (The one with more oxygen is called -ate.) • Examples: NO3- is nitrate, NO2- is nitrite. • (Exceptions: hydroxide (OH-), cyanide (CN-), peroxide (O22-).)

36. Naming Polyatomic Inorganic ions • Polyatomic anions containing oxygen with more than two members in the series are named as follows (in order of decreasing oxygen): • per-….-ate • -ate • -ite • hypo-….-ite • Examples; • ClO4- perchlorate ion, ClO3- chlorate, ClO2- chlorite, ClO- hypochlorite.

37. Road Map to Naming Monoatomic and Polyatomic Anions

38. Naming Inorganic Compounds • Polyatomic anions containing oxygen with additional hydrogens are named by adding hydrogen or bi- (one H), dihydrogen (two H), etc., to the name as follows: • CO32- is the carbonate anion • HCO3- is the hydrogen carbonate (or bicarbonate) anion. • H2PO4- is the dihydrogen phosphate anion.

39. Some Common Anions

40. Names and Formulas for Acids • The names of acids are related to the names of anions. • Acids containing anions whose names end in: • -ide becomes hydro-….-ic acid; • -ate becomes -ic acid; • -ite becomes -ous acid.

42. Names and Formulas of Binary Molecular Compounds • Binary molecular compounds have two elements. • The most metallic element is usually written first (i.e., the one to the farthest left on the periodic table). Exception: NH3. • If both elements are in the same group, the lower one is written first. • Greek prefixes are used to indicate the number of atoms.