1. … here we go again…

2. What is Matter? Is defined as anything that takes up space and has mass

3. 3 Common Characteristics Used to Describe matter: Mass the amount of matter an object has Volume the amount of space that an object takes up Size

4. States of Matter: Three States of Matter Solid : Particles in a solid are held tightly together in a rigid structure Liquid : Particles not held as tightly as in a solid (they can slip past each other-it can flow) Gas : Particles weakly attract each other – move independently at high speeds Plasma: High temperature, ionized phase of matter found on the sun

5. States of Matter

7. Definite Shape Definite Volume Indefinite Shape Indefinite Volume Solid Liquid Gas

8. Phase Changes: Solid Liquid Freezing: liquid to solid Melting: solid to liquid Melting Point (m.p.) – The temperature at which a substance melts. It is also the same temperature at which a substance freezes

9. Phase Change: Liquid Gas Evaporation/Boiling: liquid to gas Condensation: gas to liquid Boiling Point (b.p.) – The temperature at which a substance evaporates or boils. It is also the same temperature at which a substance condenses

10. Phase Change: Solid Gas Sublimation: solid to gas Deposition: gas to solid

11. Phase Changes and Volume Volume changes usually accompany changes in state.

12. Kinetic Molecular Theory All matter consists of extremely tiny particles (atoms, molecules, or ions) which are in constant motion In solids these particles are packed closely together, usually in a regular array. The particles vibrate back and forth about their average positions, but seldom does a particle in a solid squeeze past its immediate neighbors to come into contact with a new set of particles. The atoms or molecules of liquids are arranged randomly rather than in the regular patterns found in solids. Liquids and gases are fluid because the particles are not confined to specific locations and can move past one another.

13. Kinetic Molecular Theory Under normal conditions, the particles in a gas are far apart. Gas molecules move extremely rapidly because they are not constrained by their neighbors. The molecules of a gas fly about, colliding with one another and with the container walls. This random motion allows gas molecules to fill their container, so the volume of the gas sample is the volume of the container.

14. Kinetic Molecular Theory State that the higher the temperature, the faster the particles move (kinetic energy) Increasing temperature corresponds to faster and faster motions of atoms and molecules. The higher energy acts to overcome the forces of attraction… Allows solids to melt into liquids

15. Classifying Matter Why classify? Everything around us is matter… that’s a lot of stuff So we have to classify it into groups for study… and just to make it easier We’ll start big and work our way down.

16. Pure Substances vs. Mixtures If it can be separated by physical means, it is a mixture; if it can’t it’s a pure substance. Pure Substances: a sample of matter, either a single element or a single compound, that has definite chemical and physical properties 2 types of pure substances: 1. Elements 2. Compounds

17. Elements vs. Compounds If it cannot be separated by chemical means it is an element; if it can be, it is a compunds Elements Pure substances that contain only one kind of atom Can be single atoms or a molecule Things like copper, oxygen, hydrogen, helium… represented by chemical symbols on the periodic table

18. Elements Single Atoms Only one atom by itself (Monatomic) Metallic elements – arranged in a metallic lattice Non metallic – can be separate, covalently bonded, or covalently bonded in a giant lattice Molecule Usually consists of two or more atoms combined in a definite ratio. Nitrogen (N 2 )and Oxygen (O 2 ) (Diatomic) Allotrope Elements that exist in two or more different forms Ex – Oxygen vs. Ozone

19. Compounds Pure substances that are not elements Composed of more than one kind of atom Held together by chemical bonds Ionic – strong attraction between oppositely charged ions Molecular- covalent bonds holding atoms together in discrete molecules Covalent Network – atoms covalently bonded in a giant lattice

20. Mixtures A sample of matter that contains two or more pure substances Homogeneous vs. Heterogeneous Homogenous Pure substances are distributed uniformly throughout mixture Is the same throughout Often called solutions Example: Sweetened tea, 14-karat gold, gasoline, syrup, air

21. Mixtures Heterogenous Substances NOT evenly mixed Example Orange juice, tomato juice, chocolate chip cookies, salad Alloy Solid or liquid mixture of two or more metals

22. Insert chart here MATTER Can it be separated by physical means? YESNO MIXTURES PURE SUBSTANCES Is the composition uniform? Can it be decomposed by ordinary chemical means? YESNO YES Homogeneous Mixture Heterogeneous Mixture CompoundsElements

23. States of Matter The fundamental difference between states of matter is the distance between particles.

24. Intermolecular Forces The attractions between molecules are not nearly as strong as the intramolecular attractions that hold compounds together. Generally less than 15% of the values of bond energies

25. Intermolecular Forces They are, however, strong enough to control physical properties such as boiling and melting points, vapor pressures, and viscosities.

26. Intermolecular Forces Affect Many Physical Properties The strength of the attractions between particles can greatly affect the properties of a substance or solution. “Like dissolves like”

27. Phase Changes

28. Energy Changes Associated with Changes of State The heat of fusion is the energy required to change a solid at its melting point to a liquid.

29. Energy Changes Associated with Changes of State The heat of vaporization is defined as the energy required to change a liquid at its boiling point to a gas. The enthalpy change for condensation is equal but opposite in sign to the enthalpy of vaporization

30. Energy Changes Associated with Changes of State The heat added to the system at the melting and boiling points goes into pulling the molecules farther apart from each other. The temperature of the substance does not rise during a phase change.

32. Solids We can think of solids as falling into two groups: crystalline, in which particles are in highly ordered arrangement amorphous, in which there is no particular order in the arrangement of particles.

33. Attractions in Ionic Crystals In ionic crystals, ions pack themselves so as to maximize the attractions and minimize repulsions between the ions.

34. Crystalline Solids Because of the order in a crystal, we can focus on the repeating pattern of arrangement called the unit cell. The smallest repeating unit that has all of the symmetry characteristic of the way the atoms, ions, or molecules are arranged in the solid

35. Unit Cells Each unit cell is defined in terms of lattice points the points in space about which the particles are free to vibrate in a crystal. In 1850, Auguste Bravais showed that crystals could be divided into 14 unit cells, which meet the following criteria. The unit cell is the simplest repeating unit in the crystal. Opposite faces of a unit cell are parallel. The edge of the unit cell connects equivalent points.

36. Unit Cells

37. Types of Bonding in Crystalline Solids

38. Covalent-Network and Molecular Solids Diamonds are an example of a covalent-network solid, in which atoms are covalently bonded to each other. They tend to be hard and have high melting points.

39. Covalent-Network and Molecular Solids Graphite is an example of a molecular solid, in which atoms are held together with intermolecular forces They tend to be softer and have lower melting points.

40. Amorphous Solids Characteristic properties of pure crystalline solids: Melt at specific temperature Form well-defined crystals, with smooth, flat faces. Amorphous Solids don’t have these properties: Think things like glass and plastic.

41. Metallic Solids Metals are not covalently bonded, but the attractions between atoms are too strong to be intermolecular forces. In metals valence electrons are delocalized throughout the solid.