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Chapter 2 The Components of Matter He C60 Copyright © 2007 by Pearson Education, Inc. Publishing as Benjamin Cummings 1 Classification of Matter by composition Matter is anything that has mass and volume (i.e., occupies space) physically separable chemically combining 2 2.1 Elements, Compounds, & Mixtures: An Atomic Overview Pure Substance: A single chemical; one kind of matter, uniform in its chemical composition and properties. Examples: Oxygen gas, copper, sugar, water, etc… Mixture: A blend of two or more pure substances in any ratio, each retaining its identity; physical and chemical properties vary as the relative amounts of different parts change. Dissolving sugar in water creates a mixture. Can be separated into pure substances by physical changes. Evaporation and condensation can separate water from sugar. 3 A pure substance is classified as matter with a specific composition. An element when composed of one type of atom, and canNOT be broken down chemically into simpler substances. Hydrogen and oxygen are examples of elements. A compound when composed of two or more types of elements combined in a definite ratio, and can be decomposed by a chemical change into two or more other pure substances. Water is a compound composed of two parts hydrogen and one part oxygen. Water can be broken down into hydrogen and oxygen by passing an electric current through it. 4 Elements and Atoms Elements are Atoms are - Pure substances that cannot be separated into simpler substances by ordinary laboratory processes. - The building blocks of matter. - smallest identifiable units of elements. - What do atoms look like? Definitely small! If 100,000,000 copper atoms were placed side by side, they would form a line 1 cm long. Element Symbol: S, Cu, Fe,… 5 Compounds & Molecules COMPOUNDS are a combination of 2 or more elements in definite ratios. The character of each element is lost when forming a compound. MOLECULES are the smallest unit of a compound that retains the characteristics of the compound. CHEMICAL FORMULAS indicate the elements present in the compound and the relative number of atoms of each. Ethanol, C2H6O Buckyball, C60 6 2.9 Mixtures Homogeneous: A sample that has a uniform appearance and composition throughout. Solution is A homogeneous mixture. Examples: Coffee, air, brass Heterogeneous: Different phases, usually visible The composition is not uniform, it varies from one part of the mixture to another. . Examples: Carbonated beverages, salad dressings 2.2 Law of Conservation of Mass - When atoms combine into a compound, the mass of the compound is the sum of these atomic masses. - In a chemical reaction, mass is not gained or lost. the total mass of the products always equals to the total mass of the reacting substances. 8 Law of Definite Proportions All samples of the same compound contain the same proportions by mass of the component elements in a simple ratio of whole numbers. 2.3 Dalton’s Atomic Theory (1808) In Greek, ‘a’ means ‘no’; ‘tom’ means ‘cut’; ‘atom’ is the smallest, indivisible piece of matter 1) Atoms are the smallest identifiable units of elements. 2) Atoms of a given element are identical to one another, but different from atoms of any other element. 3) Atoms are rearranged in chemical reactions, but neither the number nor the types of atoms is changed in reaction 4) Compounds are formed by atoms coming together to form molecules in which the number of each type of atom is constant. What do atoms look like? Definitely small! If 100,000,000 copper atoms were placed side by side, they would form a line 1 cm long. Scanning tunneling microscope helps us to see them. 10 2.5 The Nuclear Model of the Atom -1 0 10 m Nu cleus (protons an d neu trons) Space occupied by electrons Proton Neu tron -1 5 10 m The diameter of the nucleus is about 100,000 times smaller than that of the whole atom. The correct scale would be comparable to a penny (the nucleus) in the center of the baseball field (the atom) 11 Introduction to General, Organic, and Biochemistry, 7th Ed. By Bettelheim, Brown, and March, Thomson Publishing Subatomic Particles Atom contains subatomic particles. Electrons have a negative (-) charge. Protons have a positive (+) charge. Neutrons are neutral. Like charges repel and unlike charges attract. Copyright © 2007 by Pearson Education, Inc. Publishing as Benjamin Cummings 12 Particles in the Atom 13 Atomic Number (Z) Is specific for each element. Is the same for all atoms of an element. Is equal to the number of protons in an atom. Appears above the symbol of an element. Atomic number Symbol 11 Na 14 Mass Number (A) Represents the total number of particles in the nucleus. Mass number = # of protons + # of neutrons 15 Atomic Symbol Represents a particular atom of an element by showing the element symbol Gives the mass number in the upper left corner and the atomic number in the lower left corner of the element symbol. An isotope is identified by its mass number Example: An atom of sodium with atomic number 11 and a mass number 23 has the following atomic symbols: 23 Na , 23Na, Na-23 mass number atomic number 11 16 Learning Check Write the isotope symbols for atoms with the following subatomic particles: A. 8 p+, 8n, 8e___________ B. 17p+, 20n, 17e- ___________ C. 47p+, 60 n, 47 e- ___________ 1. Which of the following pairs are isotopes of the same element? 2. In which of the following pairs do both atoms have 8 neutrons? A. 15X 15X 8 B. C. 7 12X 14X 6 6 15X 16X 7 8 17 Isotopes Isotopes are atoms of the same element that have the same number of protons, but different numbers of neutrons. 24Mg 12 25Mg 26Mg 12 12 Name of an Isotope = Elemental name–Mass number Carbon-12 = A carbon atom with 6 protons and 6 neutrons 18 Atomic Masses of the Elements How big is an atom? Very tiny!! A hydrogen atom has mass 1.67 x 10-24 g A sodium atom has mass 3.817 x 10-23 g We define a new unit: atomic mass unit (amu or u) = 1/12 the mass of a carbon-12 atom 1 amu = 1.6603 x 10-24 g So, what is the mass of a hydrogen atom, a sodium atom in amu? 11 A hydrogen atom has mass 1.008 amu A sodium atom has mass 22.99 amu Na 22.99 19 Atomic Mass (of an element): The average mass of all atoms of an element as they occur in nature. Chlorine has two natural isotopes: 75.78% is chlorine-35 at 34.968852721 amu 24.22% is chlorine-37 at 36.96590262 amu What is the atomic mass of chlorine? 0.7578 x 34.968852721 amu = 0.2422 x 36.96590262 amu = 26.50 amu 8.953 amu 35.45 amu Learning Check In nature 10B atom with a mass of 10.01294 u is 19.9% abundant, and 11B with a mass of 11.00931 u is 80.1%. Calculate atomic mass of B. The relative atomic mass of silver is 107.9 amu. If silver is composed of only Ag-107 and Ag-109, which isotope is most abundant? 21 2.6 Elements: A First Look at the Periodic Table How do we keep track of all these elements? Dmitri Mendeleev – Russian chemist (1834-1907) He listed the elements and their properties on individual cards, and organized them. Great simplification in chemistry! Tro, Nivaldo J., Chemistry in Focus: A Molecular View of our World, 3rd Ed., Thomson Brooks/Cole, 2007. 22 The Periodic Table First, Mendeleev arranged the elements in order of increasing atomic mass He examined chemical properties and found a trend! Mendeleev then arranged elements with recurring sets of properties in the same column (vertical row) and with increasing atomic number in periods (horizontal rows) He found he had 8 columns of elements that had similar chemical properties. 23 Organization of the Periodic Table Groups contain elements with similar properties in vertical columns. Periods are horizontal rows of elements in the order of increasing atomic number. Copyright © 2007 by Pearson Education, Inc. Publishing as Benjamin Cummings 24 Classifying the Elements Copyright © 2007 by Pearson Education,25 Inc. Publishing as Benjamin Cummings Properties of Metals, Nonmetals, and Metalloids Metals Are shiny and ductile. Are good conductors of heat and electricity. Nonmetals Are dull, brittle Are poor conductors of heat and electricity. Are good insulators. Metalloids Are better conductors than nonmetals, but not as good as metals. Are used as semiconductors and insulators. 26 Group Names 27 Learning Check Identify the element described by the following: A. Group 7A(17), Period 4 1) Br 2) Cl 3) Mn B. Group 2A(2), Period 3 1) beryllium 2) magnesium 3) boron C. Group 5A(15), Period 2 1) phosphorus 2) arsenic 3) nitrogen 28 ELEMENTS THAT EXIST AS POLYATOMIC or DIATOMIC MOLECULES White P4 and polymeric red phosphorus MEMORIZE THEM! S8 sulfur molecules 29 2.7 & 2.8 Compounds: Formulas, Names, & Masses Why do compounds form? - Except for the noble gases, elements usually form compounds with other elements to gain stability. Two major types of compounds: 1) ionic (metal + non-metal): loss/gain 2) molecular (non-metal + non-metal): sharing equally or unequally 30 The Formation of Ionic Compounds IONS are atoms or groups of atoms with a positive or negative charge. Taking away an electron from an atom gives a CATION with a positive charge Adding an electron to an atom gives an ANION with a negative charge. 31 Forming Cations & Anions A CATION forms when an atom loses one or more electrons. Mg --> Mg2+ + 2 e- An ANION forms when an atom gains one or more electrons F + e- --> F32 Electrostatic Forces COULOMB’S LAW As ion charge increases, the attractive force increases. As the distance between ions increases, the attractive force decreases. 33 Properties of IONIC COMPOUNDS Ionic compounds are composed of repeating positive and negative ions in large arrays (called a lattice). Formula unit of an ionic compound is the simplest whole number ratio. NaCl, salt 34 PREDICTING ION CHARGES In general, for MONOATOMIC IONS of Main Group elements metals lose electrons ---> cations ion charge = group # Mg --> Mg2+ + 2 e- nonmetals gain electrons ---> anions ion charge = group # - 8 F + e- --> F- 35 Predicting Charges on Monatomic Ions For monoatomic ions of Transition Elements: Common ion charge is 2+ or 3+ 36 Names of METAL Ions Main Groups: Na+ sodium ion Mg2+ magnesium ion Al3+ aluminum ion Transition metals Fe2+ iron(II) ion Fe3+ iron(III) ion 37 Names of NONMETAL Ions Group 4A Group 5A Group 6A Group 7A C4-,carbide N3-, nitride O2-, oxide F-, fluoride Name derived by adding -ide to stem S2-, sulfide Cl-, chloride Br-, bromide I-, iodide 38 Polyatomic Ion Is a group of atoms bonded together. Has an overall ionic charge that belongs to the whole group. In naming, simply use the name of the polyatomic ion. No ending change is needed. The most common polyatomic cation is NH4+. 39 Note: many O containing anions have names ending in –ate (or -ite). 40 Families of Oxoanions Acids Names and Their Anions Naming Ionic Compounds 1. Identify the cation and anion. 2. The cation is always named first. 3. The base name of the anion is written next, but the ending of –ide is placed on the end. 4. DO NOT use numeric prefixes in naming ionic compounds! So you really need to be sure you are naming an ionic compound. 5. Some transition metals require that you place a Roman numeral after the name to indicate the charge. 43 Naming Compounds with Polyatomic Ions The positive ion is named first followed by the name of the polyatomic ion. NaNO3 sodium nitrate K2SO4 potassium sulfate Fe(HCO3)3 iron(III) bicarbonate or iron(III) hydrogen carbonate (NH4)3PO4 ammonium phosphate 44 Naming Ionic Compounds with Variable Charge Metals Transition metals that form two different ions use a Roman numeral after the name of the metal ion to indicate the charge. 45 Writing Formulas from the Name Write a formula for potassium sulfide. STEP 1 Identify the cation and anion. potassium = K+ sulfide = S2− STEP 2 Balance the charges. K+ S2− K+ 2(1+) + 1(2-) = 0 STEP 3 Write the cation first. Add subscripts as needed 2K+ and 1S2− = K2S1 = K2S 46 Learning Check The correct formula for each of the following is: A. Copper (I) nitride E. Lead (IV) oxide B. calcium nitrate F. iron(II) hydroxide C. aluminum carbonate G. copper(II) bromide D. lithium phosphate Name the following compounds: A. Ca3(PO4)2 B. FeBr3 D. Zn(NO2)2 E. NaHCO3 C. Al2S3 47 Naming and Writing Formulas for Binary Covalent Compounds A. The first word is the name of the first element in the compound. A prefix is used to indicate the number B. The second word is the name of the second element in the compound. Its ending should be changed to –IDE. A prefix is also used to indicate the number 48 Names of Molecular Compounds Prefixes are used in the names of molecular compounds, because two nonmetals can form two or more different compounds. Examples of compounds of N and O: NO nitrogen oxide NO2 nitrogen dioxide N2O dinitrogen oxide N2O4 dinitrogen tetroxide N2O5 dinitrogen pentoxide 49 Guide to Writing Formulas STEP 1 Write the symbols in the order of the elements in the name. STEP 2 Write any prefixes as subscripts. Example: Write the formula for carbon disulfide. STEP 1 Elements are C and S STEP 2 No prefix for carbon means 1 C Prefix di = 2 Formula: CS2 50 Learning Check Name of the following molecules. CCl4 BCl3 SF6 Write the correct formula for each of the following: A. phosphorus pentachloride B. dinitrogen trioxide C. sulfur hexafluoride 51 Molecular (Formula) Masses from Chemical Formulas For a compound, the formula mass is the sum of atomic masses of the elements in the formula. Element Number of Moles Atomic Mass Total Mass in K3PO4 K 3 39.10 amu 117.3 amu P 1 31.00 amu 31.00 amu O 4 16.00 amu 64.00 amu K3PO4 52 212.3 amu Ball-and-stick model: Symbolizes atoms as balls and the electrons that connect those atoms as sticks Space-filling model: Shows the outer boundaries of the particle in three-dimensional space Representing Molecules with Formula & Models Chemical formula and Lewis diagram