Matter and Energy. Section 2.1 Chapter 2. Representations of Matter: Models and Symbols. Goal 1. Goal 2

Transcription

1 Section 2.1 Chapter 2 Matter and Energy Representations of Matter: Models and Symbols Goal 1 Goal 2 Identify and explain the difference among observations of matter at the macroscopic, microscopic, and particulate levels. Define the term model as it is used in chemistry to represent pieces of matter too small to be seen. Matter: Anything that has mass (sometimes expressed as weight) and takes up space Matter can be observed and/or thought about at different levels: Macroscopic Microscopic Particulate Macroscopic samples of matter: Mountains Rocky cliffs Huge boulders Rocks and stones Gravel Sand Macro- means large 1

2 Microscopic samples of matter: Tiny animals or plants Cells Crystals on rock surfaces Micro- means small Particulate samples of matter: Too small to see, even with the most powerful optical microscope Chemists imagine the nature of the behavior of the tiny particles that make up matter, and they use that knowledge to carry out changes from one type of macroor microscopic matter to another Macroscopic, microscopic, and particulate matter Model: A representation of something else Geologists model the earth (globe) Biologists model cells Chemists model atoms and molecules 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 Models and symbols used to represent particulate matter Ball-and-stick models Space-filling models Chemical formula and Lewis diagram 2

3 Models are represented with symbols Chemical symbols are letters that represent atoms of elements Macroscopic, particulate, and symbolic forms and representations of matter H represents an atom of hydrogen O represents an atom of oxygen H 2 O represents a molecule of water: Two hydrogen atoms and one oxygen atom Goal 3 Section 2.2 States of Matter Identify and explain the differences among gases, liquids, and solids in terms of (a) visible properties, (b) distance between particles, and (c) particle movement. States of matter: Gas The air you breathe Liquid The water you drink Solid The food you eat Kinetic Molecular Theory: All matter consists of extremely tiny particles that are in constant motion Kinetic refers to motion Molecular comes from molecule, the smallest individual particle that is present in one kind of matter Theory is a collection of general propositions that, when taken together, explain a class of related natural phenomena 3

4 The speed at which particles move is faster at higher temperatures and slower at lower temperatures There is an attraction among particles in all samples of matter The state of matter of any sample depends on temperature and the attractions among the particles that make up the sample Gas Particles are independent of one another, moving in random fashion Liquid Particles move freely among themselves, but clump together Solid Particles vibrate in fixed positions relative to one another Section 2.3 Physical and Chemical Properties and Changes Goal 4 Goal 5 Distinguish between physical and chemical properties at both the particulate level and the macroscopic level. Distinguish between physical and chemical changes at both the particulate level and the macroscopic level. 4

5 Physical Properties: Description by senses color, shape, odor, etc. Measurable properties density, boiling point, etc. Physical Changes: New form of old substance No new substances formed Charcoal is black Glass is hard The boiling point of water is 100 C Ice melts Dry ice changes to gaseous carbon dioxide A rock is ground into sand In a physical change, the particles of matter themselves are unchanged Chemical Changes: Old substances destroyed New substances formed Water decomposes to hydrogen and oxygen gases Iron rusts Food is digested When electricity is passed through certain water solutions, the water decomposes into its elements, hydrogen and oxygen. This is a chemical change. Chemical Properties: Properties defined by types of chemical changes possible Water can be decomposed to its elements Iron will rust under certain conditions Starch molecules react to form sugar molecules during digestion 5

6 Section 2.4 Pure Substances and Mixtures Goal 6 Distinguish between a pure substance and a mixture at both the macroscopic level and the particulate level. Goal 7 Pure Substance: A single chemical; one kind of matter Unique set of physical and chemical properties Distinguish between homogeneous and heterogeneous matter. Cannot be separated into parts by a physical change Mixture: A sample of matter that consists of two or more chemicals Physical and chemical properties vary as the relative amounts of different parts change Water is a pure liquid and has a constant boiling point (a physical property); The boiling point of a mixture (solution) changes as the composition of the mixture changes Can be separated into parts by physical changes 6

7 A pure substance cannot be distinguished from a mixture of uniform appearance by observation alone at the macroscopic level Solution: A homogeneous mixture Coffee, air, brass Homogeneous: A sample that has a uniform appearance and composition throughout Tea, paint, gasoline Heterogeneous: Different phases, usually visible Carbonated beverages, salad dressings Homogeneous pure substances and mixtures Section 2.5 Separation of Mixtures Goal 8 Most natural substances are mixtures; chemists separate mixtures into pure substances Nitrogen and oxygen are purified from the mixture called air Describe how distillation and filtration rely on physical changes and properties to separate components of mixtures. Pure water is purified from the mixture called natural water 7

8 Iron and sulfur form a heterogeneous mixture. Magnetism is a physical property: Iron is attracted to a magnet, sulfur is not. This physical property is the basis of the separation of this mixture Distillation: Separation of the parts of a mixture by heating a liquid until one component boils to the gaseous state. The pure gas is cooled and collected in the liquid state. Boiling is a physical change; a homogeneous mixture is changed so that at least one component is separated as a pure substance Laboratory distillation apparatus Filtration: Separation of the parts of a mixture by using a porous medium, such as filter paper, to separate components based on size Filtration is based on the physical properties of a mixture: the particle sizes of a component must be significantly larger or smaller than the pore size of the filtration medium Gravity filtration Section 2.6 Elements and Compounds 8

9 Goal 9 Goal 10 Distinguish between elements and compounds. Distinguish between elemental symbols and the formulas of chemical compounds. Goal 11 Element: Pure substance that cannot be separated into other stable pure substances Atom: Smallest unit particle of an element Distinguish between atoms and molecules. Compound: Pure substance that can be decomposed by a chemical change into two or more other pure substances The element silver and a particulate-level model of silver atoms Elements At least 88 occur in nature Examples: copper, sulfur, gold, silver 11 occur in nature as gases 2 occur as liquids (mercury and bromine) the others occur as solids Name of an element is always a single word; compound names usually two words or a multisyllabic compound word 9

10 Familiar objects that are nearly pure elements Familiar objects that are compounds Elemental symbols: Letters that symbolize elements The first letter of the name of the element, written in uppercase, is often its symbol If more than one element begins with the same letter, a second letter written in lowercase is added hydrogen, H oxygen, O carbon, C chlorine, Cl sodium, Na (from natrium) iron, Fe (from ferrum) Chemical formula: Symbolic representation of the particles of a pure substance A combination of the symbols of all the elements in the substance The formula of most elements is the same as the symbol of the element: helium, He; sodium, Na Other elements exist in nature as molecules; their formulas indicate the number of atoms of the element in the molecule: hydrogen, H 2 ; oxygen, O 2 Formula unit: A real or hypothetical particle represented by a chemical formula Ammonia molecules are real particles with the formula NH 3 : 3 atoms of hydrogen and 1 atom of nitrogen Barium chloride exists as an orderly, repeating pattern of barium and chlorine, but there is no barium chloride molecule--its hypothetical particle has the formula BaCl 2 : 2 chlorine atoms and 1 atom of barium 10

11 Law of Definite Composition (also Law of Constant Composition): Any compound is always made up of elements in the same proportion by mass (weight) The properties of a compound are different from the properties of the elements that make up the compound No matter its source, water is 11.1 parts hydrogen per 88.9 parts oxygen You are familiar with water Hydrogen and oxygen are very different from water Particulate and macroscopic views of elements and compounds Particulate and macroscopic views of elements and compounds Particulate and macroscopic views of elements and compounds Summary of the Classification System for Matter 11

12 Section 2.7 The Electrical Character of Matter Goal 12 Match electrostatic forces of attraction and repulsion with combinations of positive and negative charges. Two of the fundamental forces that govern the operation of the universe are: Force of gravity Electromagnetic force The electromagnetic force plays an important role in understanding chemistry It includes electricity and magnetism Force field: Region in space where the force is effective Electrostatic force: The force of an electrical charge that does not move A charged object exerts an invisible electrostatic force Two objects having the same charge, both positive or both negative, repel each other Two objects having unlike charges, one positive and one negative, attract each other 12

13 Matter has electrical properties There forces are responsible for the energy absorbed or released in chemical changes Goal 13 Section 2.8 Characteristics of a Chemical Change Distinguish between reactants and products in a chemical equation. Chemical Equation: Symbolic representation of the essence of a chemical change Goal 14 Reactants: Beginning substances Products: Substances formed Distinguish between exothermic and endothermic changes. C + O 2 CO 2 Reactants Products 2 H 2 O 2 H 2 + O 2 13

14 Goal 15 Exothermic Reaction: A chemical change that releases energy to its surroundings Burning charcoal: Distinguish between potential energy and kinetic energy. C + O 2 CO 2 + energy Endothermic Reaction: A chemical change that absorbs energy from its surroundings Potential Energy: Energy due to position in a field where forces of attraction and/or repulsion are present Decomposition of water: H 2 O + energy 2 H 2 + O 2 Gravitational potential energy: Position in the earth s gravitational field Electrical potential energy: Position in an electrical field Minimization of energy is one of the driving forces that cause chemical reactions to occur Kinetic Energy: Energy of motion Chemical energy comes largely from the rearrangement of charged particles in an electrostatic field The temperature of an object is related to the average kinetic energy of its particles 14

15 Section 2.9 Conservation Laws and Chemical Change The Conservation Law: In any change, the sum of mass plus energy is conserved; they are neither created nor destroyed E = m c 2 Matter is an extremely concentrated form of energy A uranium fuel pellet of this size produces energy equal to the energy that would be produced by burning about one ton of coal Goal 16 State the meaning of, or draw conclusions based on, the Law of Conservation of Mass. Law of Conservation of Mass: In a non-nuclear change, mass is conserved; it is neither created nor destroyed Goal 17 In any ordinary chemical change, Total mass of reactants = Total mass of products State the meaning of, or draw conclusions based on, the Law of Conservation of Energy. 15

16 Law of Conservation of Energy: In a non-nuclear change, energy is conserved; it is neither created nor destroyed Common events in which energy is changed from one form to another The energy lost in one form is always exactly equal to the energy gained in another form 16