Chapter 3: Matter

1. Chapter 3: Matter Properties and Change

2. States of matter • Solid: Particles are tightly packed and do not move. • Liquid: Particles are not rigidly held in place and are able to flow freely . • Gas: Particles are completely free to move. • Plasma: Plasma is ionized gas.

3. Physical Properties of Matter • A physical property of matter is a characteristic that can be observed or measured without changing the sample’s composition. • Extensive Properties: Depend on the amount of the substance being studied. • Intensive Properties: Do not depend on the amount of the substance being studied.

4. Chemical Properties of Matter • The ability of a substance to combine or change into one or more other substances is called a chemical property.

5. Changes in Matter • Physical changes: A change that alters a substance without changing its composition. • Example: Melting ice • Chemical change: A process that changes one substance in to another substance. Also called a chemical reaction.

6. Conservation of Mass • The law of conservation of mass states that matter cannot be created or destroyed. • In other words the mass of the reactants in a chemical reaction must be equal to the mass of the products.

7. Mixtures of Matter • A mixture is a combination of two or more pure substances. • A heterogeneous mixture is a mixture that does not blend smoothly throughout. • A homogeneous mixture is a mixture that does blend smoothly. • Homogeneous mixtures are also referred to as solutions.

8. Separating Mixtures • Filtration: Used to separate solids from liquids in heterogeneous mixtures • Distillation: Used to separate two liquids based on their boiling points. • Crystallization: Used to separate a pure solid from a liquid in a homogeneous mixture • Chromatography: Used to separate two liquids based on their ability to “stick” to the surface of a specific material.

9. Elements and Compounds • An element is a pure substance that cannot be separated into simpler substances by physical or chemical means. • Example: • Anything on the periodic table • A compound is made up of two or more different elements that are chemically combined.

10. The periodic table

11. More about compounds • The most important thing about compounds is that the elements that make them always combine in the same proportions. • This is called The law of definite proportions. • Percent mass:

12. Law of multiple proportions • The same two elements can combine in different proportions to create different compounds. • Example: H2O and H2O2

13. Conservation of mass • The mass reactants in a chemical reaction must be equal to the mass of the products. • Example if 10 g of Oxygen react with 10 g of carbon to form carbon dioxide they will form 20 g of carbon dioxide.

14. The atom • The atom is the smallest individual unit of a given element. • All matter is composed of atoms. • To give you an idea… • There are 6.5 x 109 people in the world. • In a single penny there are 2.9 x 1022 atoms.

15. The electron

16. Electrons • Electrons are part of all atoms. • They are very small negatively charged particles • Electrons are found out side of the nucleus (center) of the atom. • The mass of one electron is 9.1 x 10-28g.

17. The Nucleus • The nucleus of an atom is comprised of protons and neutrons. • Protons are subatomic particles that have a charge of +1 and a mass of 1 amu. • Neutrons are subatomic particles that do not have a charge and a mass of 1 amu.

18. How atoms differ • Atoms of different elements have a unique number of protons. • Example: • All atoms of carbon have 6 protons. • All atoms of nitrogen have 7 protons. • The number of protons an atom has determines what element it is.

19. Atomic Mass • The mass in grams of protons, electrons, and neutrons is so small that we do not express the mass of an atom in grams. • Instead we use Atomic Mass Units (AMU) • Since electrons are so small they do not factor into atomic mass • 1 electron = 0 AMU • Protons and neutrons are big enough to add mass to an atom. • 1 proton = 1 AMU • 1 Neutron = 1 AMU

20. Neutrons • We can calculate how many neutrons an atom has if we know what element it is as well as it’s atomic mass. • Example: • How many neutrons does an atom of carbon have if it’s atomic mass is 12 AMU • 6

21. Isotopes • Isotopes are like brothers/sisters • Isotopes are two atoms of the same element that have a different number of neutrons • Example: • Carbon – 12 • Carbon - 13

22. More About Electrons • Electrons are responsible for the chemical properties of all elements. • We can predict how an element will react once we know how many electrons it has and how they are arranged. • All electrons carry energy. • Electrons in different positions around the atom have different amounts of energy.

23. Electrons and Light • The fact that different electrons carry different amounts of energy was discovered by studying the way atoms and molecules interact with light. • Light is a form of electromagnetic radiation. • Light is made of tiny particles called photons. • These photons travel in waves.

26. Light and Energy • Just like electrons beams of light also carry energy. • The energy that electrons carry is related to the wavelength. • E = hv • Where E is energy, h is Planck’s constant, and v is frequency. • So increasing frequency is the only way to increase energy. • The only way to increase frequency is to decrease wavelength. • So light of a shorter wavelength has higher energy.

27. The Photoelectric effect • Scientists noticed that certain elements can emit energy in the form of light. • But different elements only emit light of certain colors. • i.e. certain wavelengths (AKA energy) • This led scientists to the theory that electrons are arranged in atoms according to their energy.

28. Electron Energy Levels • Electrons are arranged in levels called principle energy levels. • The principle energy levels are assigned simple numbers. • 1, 2, 3, 4 and so on. • Each principle energy level is separated into sublevels (often called sub-shells) • Each principle energy level has the same number of sub-shells as the number assigned to it. • So energy level 1 has one sub-shell • Energy level 2 has two sub-shells

29. Sub-Shells • To avoid confusion the sub-shells are assigned letters not numbers. • The letters assigned to the sub-shells are s, p, d, and f. • These sub-shells differ in the amount of electrons that they can hold. • An s sub-shell can hold 2 electrons • A p sub-shell can hold 6 electrons • A d sub-shell can hold 10 electrons • A f sub-shell can hold 14 electrons

30. Quantum theory and the atom

31. Principle energy levels • n = 1 • s • n = 2 • s • p • n = 3 • s • p • d • n = 4 • s • p • d • f

32. Electron configuration • Start by figuring out how many electrons and atom has. • Set up an electron configuration table. • Place electrons in orbitals one at a time before going to back to fill an orbital.

33. Ground State electron configurations • Low energy states are more stable that high every states. • Because of this atoms like to have their electrons in their ground states. • The aufbau principle states that each electron will occupy the lowest energy level available. • The pauli exclusion principle states that the maximum number of electrons that can occupy a single atomic orbital is two. • Hund’s rule states that single electrons with the same “Spin” must occupy each equal energy orbital before additional electrons can be added.

34. Write the ground-state electron configuration for the following elements. • Boron (B) • Strontium (Sr) • Antimony (Sb) • Rhenium (Re)

35. Alkali Metals • With the exception of hydrogen all the elements on the left of the periodic table are metals. • The elements in group 1 are called the alkali metals. • These elements are highly reactive and are usually found combined with other elements in compounds.

36. Alkaline Earth Metals • The elements in group 2 are called the alkaline earth metals. • These are also highly reactive.

37. Transition Elements • The transition elements are divided into transition metals and inner transition elements. • The inner transition elements are further divided into the actinide and lanthanide series’.

38. Nonmetals • The nonmetals make up the right hand side of the periodic table. • Nonmetals are usually found in the gaseous state or as brittle dull-looking solids.

39. Halogens • The elements in group 17 are known as the halogens. • They are also highly reactive and usually found to be part of a compound.

40. Noble Gases • The elements in group 18 are called the noble gases. • They are very unreactive.

41. The Metalliods • The elements bordering the staircase on the right hand side of the table are called the metalloids. • Metalloids have physical and chemical properties of both metals and nonmetals.

42. Organization of the Elements • The elements in the periodic table are organized by atomic number, and also by electron configuration. • The electrons in the highest principle energy level (n = 1, n = 2, n = 3….) are called the valence electrons. • All of the representative elements in the same group will have the same number of valence electrons.

44. Periodic Trends • As we move through the periodic table we see a recurring pattern of physical properties too. • One of these trends is the atomic radius of the elements. • Atomic radius increases as you go from top to bottom. • This is because the number of total electrons is increasing. • Even though the number of valence electrons is staying the same.

45. Atomic Radius

46. Ionic Radius • Ions are atoms that have either gained or lost electrons resulting in a charge on the atom. • Ions that are positively charged are called cations. • Ions that are negatively charged are called anions. • In general a cation will always be smaller than its parent atom. • In general an anion will always be larger than its parent atom.