Atoms, Isotopes, and Ions

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) • Elements are made of tiny particles called atoms. • Atoms of a given element are identical. • Atoms of different elements differ from each other in some fundamental way.

4. Dalton’s Atomic Theory (1808) • 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) • 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

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 • S2- ion _____ protons _____ electrons Name of ion: sulfide ion

30. Ion Charge and the Periodic Table

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: S2- sulfide ion

33. Naming Ions • You need to know: N3- nitride ion P3- phosphide ion O2- oxide ion S2- 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: Mg2+ and O2- • Formula:

40. Ionic Compounds Magnesium chloride • The formula is the simplest ratio of ions that have a net charge of zero. • Ions present: Mg2+ 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

44. Monoatomic anions to know

45. Practice • Name  chemical formula • Chemical formula  name