Chemistry Chapter 3 Chemical Foundations: Elements, Atoms, and Ions

1. Chemistry Chapter 3 Chemical Foundations: Elements, Atoms, and Ions

2. Chapter 3. Chemical Foundations: Elements, Atoms, and Ions Greeks- air, earth, water, and fire Robert Boyle- careful measurements and experimentation; term element 3.1 The Elements (see list of distribution by mass percent of 18 most abundant elements in earth’s crust, oceans, and atmosphere, table 3.1) (see list of elements in human body) Approx. 115 elements (88 are natural; others are artificially produced in the laboratory) Element- may be present in several different forms.

3. Top Ten

4. Table 3.1

5. Table 3.3

6. 3.2 Symbols for the Elements 1st letter always capitalized Examples: (make your own list here from what we have already studied in class) 2nd letter (if there is one) always lower case Examples: (make your own list here, too; from quiz #2 and #3) For example: CO vs. Co 3rd letter- always lower case Examples: (just a few: no quiz on these) Some symbols are based on the original Latin or Greek (or other language)(quiz #4)

7. 3.3 Daltons’ Atomic Theory (copy into notes) 1. Elements are made up of tiny particles called atoms. 2. All atoms of a given element are identical. 3.The atoms of a given element are different from those of any other element. 4. Atoms of one element can combine with atoms of other elements to form compounds. A given compound always has the same relative numbers and types of atoms. 5. Atoms are indivisible in chemical processes. That is, atoms are not created or destroyed in chemical reactions. A chemical reaction simply changes the way atoms are grouped together.

8. 3.4 Formulas of Compounds Each atom present is represented by its element symbol. The number of each type of atom is indicated by a subscript written to the right of the element symbol. When only one atom of a given type is present, the subscript is not written. Examples: SO3 (1 sulfur atom and 3 oxygen atoms) N2O5 (2 nitrogen atoms and 5 oxygen atoms) C6H12O6 (6 carbon atoms, 12 hydrogen atoms, 6 oxygen atoms) Problem: a molecule w/ 1 uranium atom and 6 fluorine atoms (UF6) A molecule w/ 1 aluminum atom and 3 chlorine atoms (AlCl3)

9. Figure 3.2: Representation of NO, NO2, and N2O.

10. 3.5 The Structure of the Atom John Dalton- modern atomic theory J.J. Thomson experiment- cathode ray tube discovery- electron model- “plum pudding” Ernst Rutherford experiment- gold foil discovery- nucleus model- nuclear James Chadwick (and Rutherford) discovery- neutron model- more detailed nuclear model

11. Figure 3.7: Schematic of a cathode ray tube.

12. Figure 3.3: Plum Pudding model of an atom.

13. Figure 3.5: Rutherford’s experiment.

14. Figure 3.6: Results of foil experiment if Plum Pudding model had been correct.

15. Figure 3.6: Actual results.

16. 3.6 Introduction to the Modern Concept of Atomic Structure Particle relative mass relative charge Mass, in amu Electron 1 1- 0 Proton 1836 1+ 1 Neutron 1839 none 1 Elements of different atoms differ in the number of protons. 1 proton – Hydrogen 4 protons- Beryllium 2 protons- Helium 5 protons- Boron 3 protons- Lithium 6 protons- Carbon

17. Figure 3.9:A nuclear atom viewed in cross section.

18. 3.7 Isotopes Atomic Number (symbol Z) - # of protons in the nucleus of an atom Mass number (symbol A) - sum of # of protons and # of neutrons in an atom Isotope- atoms of an element with the same # of protons, but different # of neutrons. Symbol Identity AZX where X is symbol of element Examples: 9038Sr (# protons = 38 = # electrons; 90-38 = 52 = # neutrons) 20180Hg (# protons = 80 = # electrons; 201-80 = 121 = # neutrons) Problem: Write the symbol identity for silver (Z = 47) w/ 61 neutrons (Answer: 10847Ag) Write the symbol identity for phosphorus (Z = 15) w/ 17 neutrons (Answer: 3215P)

19. Figure 3.10: Two isotopes of sodium.

20. (not in the textbook) Dobereiner organized elements of similar properties into “triads.” He also grouped them with attention to the difference in atomic mass. This was one of the early attempts to organize the random information about elements. Li Cl Mg Na Br Ca K I Sr

21. (not in the textbook) Newlands “Law of Octaves” Newlands, an Australian scientist, organized elements according to increasing atomic mass and noticed that elements lined up with similar properties. He observed that the elements 1 & 8 were similar, 2 & 9 were similar, and so on. Other scientists were critical of his work because he used a music term to describe a scientific observation. Li Be B C N O F Na Mg Al Si P S Cl What group of elements was not included in his analysis? Why ?

22. 3.8 Introduction to the Periodic Table Dmitri Mendeleev Henry Mosely (not in your textbook; learn about him!) Groups- a vertical column in the periodic table Families- a group of elements with similar properties Alkali Metals- Group 1; extremely reactive; low density; shiny Alkaline Earth Metals- Group 2; very reactive; slightly higher density and hardness than Group 1. Halogens- Group 17; nonmetals, extremely reactive; colorful Noble Gases- Group 18; inert gases Transition Metals- Groups 3-12; various colors (most of the elements that we think of as metals) Periods- a horizontal row in the periodic table; there are 7 periods

23. Figure 3.11: The periodic table

24. Metals- solids at room temperature (except Mercury); malleable (able to be bent); ductile (able to be pulled into a fine wire); shiny (luster); good conductor of heat and electricityNonmetals- many are gases, Bromine is a liquid at room temperature; not malleable, not ductile; not shiny; poor conductorSemimetals (metalloids)- properties similar to both metals and nonmetals. Si- shiny, high melting pt., poor conductor of electricity (compared to most metals), but can conduct electricity at temperatures where most metals would have melted.

25. Figure 3.12: Elements classified as metals and nonmetals.

26. 3.9 Natural States of the Elements Noble Gases- gases which do not combine chemically with other elements (He, Ne, Ar, Kr, Xe, Rn) Noble Metals- metals which are generally chemically unreactive (Au, Pt, Ag) Diatomic Molecules: H2, O2, N2, F2, Cl2, Br2, I2 (BrINClHOF) LEARN THIS LIST! Elemental Liquids- Br and Hg Allotropes- different forms of the same element Example: C (carbon)- graphite, diamond, buckyballs)

27. Figure 3.13: A collection of argon atoms.

28. Figure 3.14: Nitrogen gas contains NXN molecules.

29. Table 3.5

30. 3.10 Ions Ion- an atom which has lost or gained electrons Cation- an ion with a positive charge; lost electron(s) Examples: lithium ion Li+ Sodium ion Na+ Magnesium ion Mg2+ Anion- an ion with a negative charge; gained electron(s) Examples: fluoride ion F- Oxide ion O2- Nitride ion N3- (note the difference in naming cations and anions!)

31. Figure 3.19: The ions formed by selected members of groups 1, 2, 3, 6, and 7.

32. 3.11 Compounds That Contain Ions Criss cross method of writing formulas for ionic compounds: Na+ and Cl- NaCl Mg2+ and Cl-  MgCl2 Li+ and N3-  Li3N Reduce subscripts to lowest whole number ratio. Always write cation first, anion second. Examples: Ca2+ and Cl- CaCl2 Na+ and S2-  Na2S Ca2+ and P3-  Ca3P2 Problems: K+ and I-  KI Mg2+ and N3-  Mg3N2 Al3+ and O2-  Al2O3

33. Figure 3.20: Pure water does not conduct a current.

34. Figure 3.20: Water containing dissolved salt conducts a current.