Atomic Structure Ch. 3

Early Philosophers Democritus – 400 BC Aristotle Thought matter was made of individual, indivisible particles – “atomus” Aristotle Disagreed with Democritus Thought matter was continuous His idea was accepted until late 1700s

Beginnings of Modern Atomic Theory 1790s – Quantitative analysis begins with improved measuring equipment. Chemists studied chemical reactions Law of Conservation of Mass – Matter cannot be created or destroyed Law of Multiple Proportions – different compounds with the same elements will have different ratios of those elements (CO and CO2) Law of Definite Proportions – the same compound will have the same proportions by mass no matter what the sample size

Dalton’s Theory All matter is composed of atoms. Atoms of the same element are identical. Atoms cannot be divided, created or destroyed. Atoms can combine in simple whole number ratios to form compounds. Atoms are rearranged, separated, or combined in chemical reactions.

Modern Atomic Theory Dalton helped turn Democritus’s idea into scientific theory. All matter is composed of atoms. Atoms of one element differ in properties from atoms of another element.

Discovery of the Electron JJ Thomson used cathode ray tubes to study electricity and matter

Observations of the Cathode Ray Experiments An object placed in the tube caused a shadow A paddle wheel between the electrodes moved Cathode rays were deflected with the negative end of a magnet and attracted by the positive end of a magnet Video on Cathode Ray Tube

Conclusions of the Cathode Ray Experiments The particles in the cathode ray have mass. The particles are negatively charged. The particles are very small. Called Electrons

Discovery of the Nucleus Rutherford used the Gold Foil Experiment to study the atom Alpha particles are fired at the thin gold sheet. Animation of the experiment

Observations of the Gold Foil Experiment Some particles were deflected by the gold foil. Some at sizeable angles. Most particles passed straight through the gold foil.

Conclusions of the Gold Foil Experiment The nucleus is very small. The nucleus is very dense. The nucleus is positively charged. Video of Rutherford’s experiment

Subatomic Particles Protons (p+) Neutrons (n0) Electrons (e-) Positively charged, found in nucleus Mass number of 1 Neutrons (n0) Neutral, found in nucleus Electrons (e-) Negative, found in electron cloud surrounding nucleus, equals the number of protons in a neutral atom Mass number of 0, Negligible mass

“Picture” of cesium and gallium atoms created using a scanning tunneling microscope

Subatomic Particles 3 quarks = 1 proton or 1 neutron Quarks He component of protons & neutrons 6 types He 3 quarks = 1 proton or 1 neutron

About Quarks… Protons and neutrons are NOT fundamental particles. Protons are made of two “up” quarks and one “down” quark. Neutrons are made of one “up” quark and two “down” quarks. Quarks are held together by “gluons”

Atomic Number Equals the number of protons Identifies the element Found as the whole number on the Periodic Table EX: Carbon has the atomic number of 6, so carbon has 6 protons in the nucleus.

Mass Number Equals the number of protons plus neutrons NOT the mass found on the Periodic Table Has to be provided to you by either Hyphen Notation Nuclear Notation

Nuclear Symbol Notation Mass Number (p+ + n0) Atomic Number (p+ only, it is not necessary to write the atomic number) How many protons does this element have? 6 How many neutrons does this element have? 8 How many electrons does this element have? 6

Hyphen Notation Chlorine – 35 or Cl – 35 How many protons? 17 How many neutrons? 18 How many electrons? 17 Mass Number

Relative Atomic Mass Masses of atoms are very small Ex: One O atom = 2.657 x 10-23 g Carbon-12 has been assigned 12 atomic mass units This is the standard for comparison 1 amu = 1/12th the mass of Carbon-12

C. Relative Atomic Mass atomic mass unit (amu) 12C atom = 1.992 × 10-23 g atomic mass unit (amu) 1 amu = 1/12 the mass of a 12C atom 1 p = 1.007276 amu 1 n = 1.008665 amu 1 e- = 0.0005486 amu © Addison-Wesley Publishing Company, Inc.

Isotopes Elements with the same number of protons but different number of neutrons. Have different mass numbers. Ex: Carbon-12 and Carbon-14 Both have 6 protons, but carbon-12 has 6 neutrons and carbon-14 has 8 neutrons.

Another Example of an Isotope Nucleus Hydrogen–1 (protium) 1 Hydrogen-2 (deuterium) Hydrogen-3 (tritium) 2

Average Atomic Mass Average mass of the isotopes found in nature Mass found on Periodic table Based on the percent abundance in nature Example: 99.985% of hydrogen found in nature is hydrogen-1 and 0.015% is hydrogen-2. This creates an average mass of 1.00794 amu.)

Example of Average Atomic Mass Isotope Symbol Composition of the nucleus % in nature Carbon-12 12C 6 protons 6 neutrons 98.89% Carbon-13 13C 7 neutrons 1.11% Carbon-14 14C 8 neutrons <0.01%

D. Average Atomic Mass Avg. Atomic Mass weighted average of all isotopes on the Periodic Table round to 2 decimal places Avg. Atomic Mass

D. Average Atomic Mass EX: Calculate the avg. atomic mass of oxygen if its abundance in nature is 99.76% 16O, 0.04% 17O, and 0.20% 18O. Avg. Atomic Mass 16.00 amu

D. Average Atomic Mass EX: Find chlorine’s average atomic mass if approximately 8 of every 10 atoms are chlorine-35 and 2 are chlorine-37. Avg. Atomic Mass 35.40 amu

The Mole Chemist’s counting number SI unit for amount of a substance Avogadro’s Number = 6.022 x 1023 There are always 6.022 x 1023 particles in one mole of any substance Ex: 1 mole of Carbon atoms = 6.022 x 1023 atoms Chemists use a really big number because atoms are so small.

Molar Mass Mass of one mole of any substance Relates the amount of atoms to their mass Helps chemists know how many atoms are present – Remember atoms are too small to count!!! Chemists can convert between mass (grams) and moles using the Periodic Table

Mass (Gram) to Mole Conversions Molar mass is the conversion factor used to convert grams to moles Found for every element on the Periodic Table – Round to two decimal places!!!!! What is the molar mass of carbon? 12.01 g/ 1 mol What is the molar mass of chlorine? 35.45 g/ 1 mol

Example of a Conversion How many grams of lithium are in 3.50 moles of lithium? Look up the molar mass on the PT. Write down given 3.50 mol Li 6.94 g Li = g Li 45.1 1 mol Li Multiply. Round to correct sig figs. Don’t forget the unit! Determine which conversion factor is necessary. Place the unit to cancel on the bottom.

Conversion Practice Use dimensional analysis to solve these problems. How many moles of carbon are present in 11.9 g of carbon? What is the mass of 3.50 moles of copper? How many atoms are present in 2.75 moles of lithium? How many atoms are present in 95.5 grams of barium?