Chapter 3

1. Chapter 3 • Elements • Periodic Table • Compounds

2. Elements • All words in the English language are the result of the 26 letters of the alphabet. • All substances on our planet and the universe are the result of about 88 known “building blocks” called the elements. • Why 88?

3. Elements • An element cannot be decomposed into a simpler substance. • The smallest particle of an element is the atom. • Atoms are incredibly small.

4. Elements • Every element has a one or two letter designation called its _____________. • In writing a symbol, the first letter is capitalized and the second letter is lowercase. • C = carbon • Co = cobalt • But, CO = a compound of carbon and oxygen!

5. Distribution of the Elements • Most elements are solids at 25oC. • Br and Hg are liquids • H, N, O, F, Cl, He, Ne, Ar, Kr, Xe, and Ra are gases • Silicon and oxygen make up almost 75% of the Earth’s crust.

6. Distribution

7. History of the Elements • Some elements were known to ancient man like Copper and Gold. • Names and symbols are based in Latin language • Copper = Cuprum; symbol Cu • Gold = Aurum; symbol Au • www.ptable.com

8. Common Elements

9. Periodic Table • Dimitri Mendeleev, 1869 • Elements are arranged into vertical columns called groups. • Elements within the same group tend to react in the same way. • Horizontal row = period.

10. Periodic Table • Main (representative) groups = A • Transition groups = B • Groups with special names • Alkali metal, alkaline earth metal, halogens, and the noble gases. • Actinide and lathanides.

11. Periodic Table • A stair-step line on the right-hand side of the chart separates the metals from the non-metals. • Metals are solids, have luster (shine), are good conductors, are malleable,and are ductile. Most have a high density and high melting point.

12. Periodic Table • Non-metals are dull, are poor conductors, and have low melting points and densities. • Metalloids are elements that fall on or near the line and have properties of both metals and non-metals. • B, Si, Ge, As, Sb, Te, and Po • Hydrogen – a special exception.

13. Groups

14. Learning Check • Describe each in terms of a Group, Period, and Metal, non-metal, or metalloid.

15. Elements in Nature • Elements are rarely found in their pure form in nature. • Exceptions are metals like Gold, Silver, and Platinum; non-metals like Nitrogen and Oxygen; and the noble gases. • Most are combined with other elements as compounds.

16. Diatomics • Some important elements occur as two atoms bonded together. • The list is:

17. Compounds • Compounds are classified as either Molecular or Ionic. • Molecular (aka covalent) compounds are formed between two or more non-metals who share electrons. • Ionic compounds are formed between a cation (+) and an anion (-).

18. Compounds • Metals lose electrons and form cations. • Na  Na+1 + 1e- • Non-metals gain electrons and form anions. • Cl + 1e-  Cl-1 • Thus, an ionic compound will have a metal and a non-metal in its formula.

19. Compounds • Law of Constant Composition – any sample of a compound will always have the same ratio or mass percentage of its elements. • Water always has two hydrogen atoms and one oxygen atom per formula. • Sodium chloride always has one sodium and one chlorine atom per formula.

20. Chemical Formulas • Are a shorthand method for writing names. • Subscripts indicate the number of each of the elements in the compound. • A subscript of one is NEVER written. • Ex) CO2, MgBr2, C6H12O6.

21. Learning Check • Decide whether the following would be a molecular or ionic compound.

22. Chapter 5 • Atomic Theory • Discovery of Atomic Structure • Subatomic Particles • Nuclear Atom

23. Dalton’s Model • All matter is made up of tiny particles called atoms. • All atoms of a given element are similar to one another; atoms of different elements are different from each other. • Atoms of two or more different elements combine to form compounds. • A chemical reaction involves the rearrangement of atoms into new combinations. Atoms are never created nor destroyed in a chemical reaction.

24. Parts of an Atom • Experiments performed around the turn of the previous century (~1900), showed that atoms were made of several types of particles – collectively referred to as subatomic particles. • These experiments showed that three types of particles were present in an atom.

25. Subatomic Particles • A proton has a +1 charge and an approximate mass of 1 amu. • Note: an atomic mass unit (amu) is equal to 1/12 of the mass of a Carbon atom with 6 protons and 6 neutrons. • A neutron has no charge, but does have a mass of about 1 amu. • An electron has a –1 charge and a mass so small that we usually say that it weighs 0 amu.

26. Atomic Structure • Ernest Rutherford performed an experiment called the “Gold Foil” experiment in 1911. • He used an alpha particle (2P + 2N) source and fired them at a piece of very thin gold foil. • He expected all of the particles to pass straight through. However, some were deflected and some were even reflected backwards. • In Rutherford’s words, it was as if he had shot a cannonball at a piece of tissue paper and have it bounce backwards.

27. Gold Foil Experiment

28. Results • Only 1 in 8000 alpha particles is scattered. • Scattering occurs when an alpha particle encounters a gold nuclei. • A nucleus is very small and contains both the protons and the neutrons. Thus, it contains almost all of the mass of an atom. • This very dense center is surrounded by the electron cloud, which is occupied by the fast moving electrons. • Thus, an atom is MAINLY EMPTY SPACE.

29. Atomic Model

30. Atomic Number & Mass Number • All atoms of the same element have the same number of protons. • This distinguishes one element from another. • The number of protons is also called the atomic number. • This is always the integer found on the periodic chart with each chemical symbol.

31. Atomic Number & Mass Number • Atoms are electrically neutral. Thus, each element must have an equal number of protons and electrons. • The mass number of an atom is equal to the sum total of the protons and neutrons in the nucleus. • Mass number and atomic weight (found on the periodic chart) are NOT the same thing.

32. Study Check

33. Isotopes • All atoms of one element have the same number of protons. • But, they can have different numbers of neutrons, and hence, a different mass number. • These different versions of atoms from one element are called isotopes.

34. Isotopes • Use the chemical symbol, atomic number (Z), and mass number (A) as seen below. • Can also list symbol followed by mass number.

35. Atomic Mass • The masses found for each element on the periodic chart are the weighted average of all the known isotopes for that element. Example: Chlorine has only two known isotopes – Cl-35 and Cl-37. Cl-35 is found 75.5% of the time and Cl-37 is found 24.5% of the time.

36. Isotope Mass X Percent = Contribution (approximate) to total mass 35 amu X 0.755 = 26.4 amu 37 amu X 0.245 = 9.1 amu Totals 1.000 35.5 amu • With all elements, round A.W.’s to one decimal place.

37. Study Check • What is the Atomic Weight of each element rounded to 0.1amu? • Na • Si • Cl • K