1. Atoms, Isotopes, and Ions 4.3 – 4.7, 4.10, 4.11

2. Atomic Theory In 1808 John Dalton proposed atomic theory. Dalton’s theory explained several laws known at the time. –Law of conservation of matter –Law of definite proportions –Law of multiple proportions

3. Dalton’s Atomic Theory (1808) 1.Elements are made of tiny particles called atoms. 2.Atoms of a given element are identical. 3.Atoms of different elements differ from each other in some fundamental way.

4. Dalton’s Atomic Theory (1808) 4.Atoms of one element can join with atoms of other elements to form compounds. A given compound is always made of the same elements combined in the same ways.  Explains the law of multiple proportions and the law of definite composition.

5. Dalton’s Atomic Theory (1808) 5.Atoms are indivisible in chemical reactions. Chemical reactions change how atoms are grouped (bonded) together.  Explains the law of conservation of matter.

6. Atomic Theory Dalton’s proposal lead to much research as to the nature of the atom. In the late 1800’s chemists/physicists determined that the atom is made up of smaller, subatomic, particles.

7. Atomic Theory - 1910 ~1896, JJ Thomson demonstrated that atoms can emit negative particles. –Called these particles electrons. –Since atoms are neutral he also proposed that they must contain positive particles. These + particles were not fully described/named until 1919.

8. Atomic Theory - 1910 ~1910 Lord Kelvin proposed the “plum pudding” model of the atom. –Proposed that electrons were scattered within a “cloud”/pudding of positive charge.

9. Atomic Theory - 1911 ~1911 an experiment was conducted in Ernest Rutherford’s lab that showed the “plum pudding” model to be incorrect. –Experiment was conducted by Geiger and Marsden and the findings interpreted by Rutherford. –See page 84

10. The gold foil experiment What they did – see board What they found – see board What Rutherford concluded.

11. Rutherford’s Model of the Atom First to propose a nuclear atom. Rutherford proposed that: –the atom must have nearly all its mass, and positive charge, in a central nucleus about 10,000 times smaller than the atom itself. –Most of the atom is empty space and the electrons are scattered through out this empty space.

12. A New Model of the Atom Expected based on Plum pudding model Rutherford’s model Based on ”his” results

13. Subatomic Particles Rutherford continued to study the atom and the positive matter of the atom. –1919, + particle named the proton ~1932 James Chadwick proposed the existence of a third subatomic particle, the neutron.

14. Subatomic Particles Subatomic Particle ChargeMass, amuLocation in atom Electron (e - ) 0 amuOutside of nucleus Proton (p)+1~1 amuNucleus Neutron (n)0~1 amuNucleus

15. Mass of Subatomic Particles Protons and neutrons have ~ the same mass (in the range of 10 -24 g). –Neutrons are slightly heavier. –Mass is expressed in amu Atomic mass unit (amu) – 1/12 the mass of a carbon-12 atom

16. Mass of Subatomic Particles The mass of the electron is tiny as compared to that of the proton and neutron. –Therefore, the electron’s mass is considered to be ~0 amu when calculating the mass of an atom.

17. Subatomic Particles and the Elements Each element has a unique number of protons. –Number of protons defines the element.

18. Subatomic Particles and the Elements Since atoms are neutral, for every proton there is a/n _________. When atoms interact to form compounds, it is their ___________ that “intermingle”.

19. Terms Atomic number = number of protons in an atom –Also indicates the number of electrons in the atom. –Finding atomic number on the periodic table.

20. Terms Mass number = sum of the # of protons and the # neutrons in the nucleus of an atom –FOR MOST ELEMENTS THE MASS NUMBER IF NOT ON THE PERIODIC TABLE. You will be given enough information to determine mass number or number of neutrons.

21. Terms Isotopes = atoms of a given element that differ in mass number –Isotopes have the same number of _____________. –Isotopes differ in the number of _______.

22. Isotopes Writing atomic symbols for isotopes –See board and pg 87

23. FAQ - Isotopes When is mass number found on the periodic table? What’s the atomic mass? Is it the same as the mass number?

24. Practice Start # 42 on page 110.

25. Ion Formation Ions are formed when atoms gain or lose electrons. –Proton and neutron number are unchanged when an ion forms.

26. Ions - Terms Ion – charged atom or group of atoms –Cation = positively charged ion Metals form cations. –Anion = negatively charged ion Nonmetals form anions.

27. Ions Na atom _____ protons_____ electrons Na + ion _____ protons_____ electrons Name of ion: sodium ion

28. Ions Calcium atom _____ protons_____ electrons Ca 2+ ion _____ protons_____ electrons Name of ion: calcium ion

29. Ions Sulfur atom _____ protons_____ electrons S 2- ion _____ protons_____ electrons Name of ion: sulfide ion

30. Ion Charge and the Periodic Table Group #Ion Charge# e - lost/gained I A1+1 e - lost II A2+2 e - lost III A metals3+3 e - lost V A nonmetals3-3 e - gained VI A nonmetals2-2 e - gained VII A1-1 e - gained

31. Naming Ions Name of a monatomic cation is the name of the element –Examples: Ca 2+ calcium ion Al 3+ aluminum ion K +

32. Naming Ions Monatomic anions are named by changing end of the name of the element to “ide” Example: S 2- sulfide ion

33. Naming Ions You need to know: N 3- nitride ion P 3- phosphide ion O 2- oxide ion S 2- sulfide ion F - fluoride ion Cl - chloride ion Br - bromide ion I - iodide ion

34. Ionic Compounds Structure –In an ionic compound there is a regular arrangement of oppositely charged particles. –Ions are arranged in a 3-D crystalline structure that maximizes attractive forces and minimizes repulsive forces. Also called a lattice structure See page 102

35. Ionic Compounds Physical Properties – all are related to the structure of the compounds –Solids at room temperature –Relatively high melting and boiling points –No vapor pressure Meaning… they don’t evaporate –Electrolytes Conduct electricity when melted or dissolved in water

36. Ionic Compounds The chemical formula for an ionic compound represents the lowest, whole number ratio of the component ions that has a net charge of zero. Total positive charge = total negative charge

37. Ionic Compounds Name the compound by naming the ions.

38. Ionic Compounds Writing formulas for and naming binary ionic compounds –Magnesium oxide

39. Ionic Compounds Magnesium oxide –The formula is the simplest ratio of ions that have a net charge of zero. –Ions present: Mg 2+ and O 2- –Formula:

40. Ionic Compounds Magnesium chloride –The formula is the simplest ratio of ions that have a net charge of zero. –Ions present: Mg 2+ and _____ –Formula:

41. Ionic Compounds Practice –Note we are currently applying the content of 4.11 and 5.2 (type I binary ionic compounds)

42. Types I Binary Compounds Compound between a metal and a nonmetal –Metal forms only one ion Name the cation and then the anion. –Name of the cation is the name of the element –Name of the anion is the name of the nonmetal with the ending changed to “ide”

43. Monoatomic cations to know Group #Charge on ionexamples IA+1Na 1+ sodium (ion) K 1+ potassium (ion) IIA+2Mg 2+ magnesium (ion) IIIA metals +3Al 3+ aluminum (ion)

44. Monoatomic anions to know Group #Charge on ionexamples VA-3N 3- nitride (ion) P 3- phosphide (ion) VIA-2O 2- oxide (ion) S 2- sulfide VIIAF 1- fluoride (ion) Cl 1- chloride (ion) Br 1- bromide (ion) I 1- iodide (ion)

45. Practice Name  chemical formula Chemical formula  name