Chapter 2- Atoms

1. Chapter 2- Atoms HW-2.10, 2.11, 2.15, 2.17, 2.25, 2.29, 2.45, 2.48, 2.53, 2.54, 2.66, 2.79

2. Ancient Thinkers • Pondered what matter was made of • Democritus and followers believed matter was composed of tiny particles • Called particles atoms • Said their were different types of atoms that add different properties • Zeno of Elea believed matter was infinitely divisible • Atoms- the basic unit of matter

3. Classifications of Matter

4. Definitons • Elements- a substance that consists of identical atoms • Compounds- a pure substance made up of two or more elements in a fixed ratio by mass. • Formula- gives the ratio of a compounds constituent elements and identifies each element by its atomic symbol • Mixture- a combination of two or more pure substances

5. Types of Mixtures • Homogeneous mixture- composition the same throughout • Example: Air • 78% Nitrogen 21% Oxygen • Heterogeneous mixture- nonuniform composition • Example: Blood • Note: Mixtures consists of two or more pure substances with different physical properties. This properties can be used to separate the mixture.

6. Dalton’s Atomic Theory • All matter is made up of very tiny, indivisible particles called atoms • All atoms of the same element have the same chemical properties • Compounds are formed by the chemical combination of two or more different kinds of atoms • A molecule is a tightly bound combination of two or more atoms that acts as a single unit

7. Laws • Law of Conservation of Mass • Matter can be neither created nor destroyed (This does not contradict Dalton!!) • Law of Constant Composition • Compounds are always made up of elements in the same proportion by mass

8. Types of Elements • Monoatomic Elements- elements that consist of single atoms not connected to each other. Ex. He, Ne • Diatomic Elements- elements that exist as pairs of atoms, bonded together. Ex. O2 Cl2 Br2 I2 H2 N2 • Polyatomic Elements- elements that exist with more than two atoms bonded together. Ex. O3 P4 S8 Diamond

9. Subatomic Particles • Proton- subatomic particle found in the nucleus of an atom and has a positive charge • Electron- subatomic particle found in the space surrounding the nucleus and has a negative charge • Neutron- subatomic particle found in the nucleus and has no charge

10. Atom Properties • Mass Number- the sum of the number of protons plus the number of neutrons • Atomic Number- the number of protons in the nucleus of an element • Symbol for an atomic nucleus- The element symbol with the mass number as a superscript on the left and the atomic number as a subscript on the left.

11. Properties of Atoms (cont.) • Isotopes-atoms with the same number of protons but different numbers of neutrons • Properties of isotopes of the same element are almost identical except for radioactivity properties. • Atomic weight- of an element in the Periodic Table is a weighted average of the masses (in amu) of its isotopes found on Earth.

12. Periodic Table • First created by Mendeleev • Still used today • Started by arranging elements by weight beginning with hydrogen • Noticed trends so he started arranging them in rows (periods) and started a new period every time he found an element with properties like hydrogen. • Discovered that elements in other groups had similar properties as well.

13. Definitions • Main-group elements- elements in groups 1A, 2A, and 3A-8A. • Transition elements- elements in B groups • Inner transition elements- elements 58-71 and 90-103

14. Classes of Elements • There are three classes • Metals- most elements in the Table. They are solid (except Hg), shiny, conductors, ductile (wires), and malleable (sheets) • Alloys- solutions of one or more metals dissolved in another metal. Examples: Bronze (copper and tin) Pewter (tin, antimony, and lead)

15. Classes (cont) • Nonmetals- with the exception of Hydrogen, they are on the right side of the table. • Non conductive (exception graphite) • Tend to accept electrons in reactions • Metalloids- B, Si, germanium, arsenic, antimony, and tellurium • Have some properties of both

16. Trends in the Periodic Table • Sizes increases as you move down a column • For the halogens, group 7A, MP and BP increase as you move top to bottom • Ionization energy- is a measure of how difficult it is to remove an electron from an atom in the gaseous state • Ionization energy increases as you move up a group and from left to right across a period

17. Electron Configuration • Electrons are located in the considerably larger space outside the nucleus. • The lowest possible energy level an electron can be in is the Ground State. • The energy of electrons is quantized and there are only certain energy levels an electron can be in.

18. Electron Configuration • Electrons in atoms are confined to specific regions of space called Principal energy levels or more simply, shells. • Shells are numbered 1,2,3,4, and so on • Electrons in the first shell are closest to the nucleus, are held most strongly, and therefore are the lowest in energy. • As you move away, energy increases

19. Electron Configuration • Shells are divided into subshells designated by the letters s, p, d, and f. • Within subshells electrons are grouped into orbitals. • Orbital- is a region of space that can hold two electrons. • The first shell has one orbital called the 1s • The second shell contains one s orbital and three p orbitals.

20. Electron Configuration • P orbitals come in sets of three and holds a total of six electrons • The 3rd shell contains one s orbital, three p orbitals, and five d orbitals.

21. Orbital Shapes • All s orbitals have the shape of a sphere with the nucleus in the middle. • Each 2p orbital has the shape of a dumbbell with the nucleus at the midpoint of the dumbell. • The three 2p orbitals are at right angles to each other.

22. Electron Configuration • Electron Configuration- a description of the orbitals that its electrons occupy • In the ground state of an atom only the lowest energy orbitals are occupied, all others are empty • We determine the ground state configuration using the following rules:

23. Rules • Orbitals fill in the order of increasing energy from lowest to highest • Each orbital can hold up to two electrons with spins paired • When there is a set of orbitals of equal energy, each orbital becomes half-filled before any of them becomes completely filled.

24. 3 types of configuration notation • Orbital box diagrams- use boxes or lines to represent an orbital • Condensed configuration- write only shell number, orbital, and number of electrons in that orbital • Noble Gas Notation- use noble gas symbol for inner shell designation

25. Showing electron config. • Usually only show the outer shell electrons • Outer-shell electrons are called valence electrons • The energy level in which valence electrons are found is called the valence shell. • To show the outmost electrons of an atom, we commonly use a representation called a Lewis Dot Structure

26. Lewis Dot Structures • Lewis Dot Structures- shows the symbol of the element surrounded by a number of dots equal to the number of electrons in the outer shell of an atom of that element. • Valence electrons- are electrons in the outermost shell. • Examples-

27. Another trend in the Periodic Table • Elements in a group have the same electron configuration in their outer shell!! • That mean the Lewis Structure will be the same as well, except for the atomic symbol.

28. Atoms vs. Ions • Atoms are neutral species that have the same number of electrons and protons. • During typical chemical reactions, the number of protons and neutrons does not change • The number of electrons does change!! Atoms can gain or lose electrons. • Ions are a compound or element that has gained or lost an electron and is now charged!!

29. Atoms vs. Ions • Example- • Ionization Energy- a measure of how difficult it is to remove an electron from an atom. • Ionization Energy increases as you move up a group and from left to right on the periodic table.