Atomic Structure Atom = the smallest particle of an element that retains the properties of that element

Law of Conservation of Mass Matter cannot be created or destroyed What happened in Lab 2?

Law of Definite Proportions A chemical compound contains the same elements in exactly the same proportions by mass regardless of the size or source of the sample H2O is water, all water is H2O, or it’s NOT water H2O2 is hydrogen peroxide, NOT WATER

2 regions of the atom Nucleus Electron cloud Protons (p+) Neutrons (n0) Electron cloud Electrons (e-)

Discovery of Electrons Cathode ray experiments Thomson e- have large charge:mass ratio Mass of e- was determined by measuring charge and calculating mass based on ratio Mass of e- = 9.109 x 10-31

Atoms are electrically neutral Negatively charged e- are balanced by positively charged particles (p+) Mass of e- is negligible, other particles must account for atomic mass Rutherford’s model = plum pudding model Rutherford, Geiger, and Marsden = densely packed, positively-charged nucleus

Subatomic Particles With 1 exception, all atomic nuclei contain p+ and n0 Charge of p+ = charge of e- n0 has NO charge Atom is neutral, because #p+ = #e- Mass of p+ > mass of e-

Forces in Nucleus + attracts – + repels +, so how are the p+ in the nucleus held together? When p+ are in close proximity to other p+ and n0, strong forces are created = nuclear forces But a nucleus can only hold so many protons in proximity and remain stable

Atomic Number and Mass Number Atomic # = # p+ #p+ = #e- Mass number = #p+ + #n0 Remember: mass of e- is negligible The atomic number or # of p+ determines the atom’s identity (which element) The atomic number for every element can be found in the Periodic Table Element builder

Atomic # appears above symbol Atomic mass appears below symbol Why is it a decimal?

Average Atomic Mass The mass number on a PT is the average atomic mass Isotopes = atoms of the same element that have different numbers of neutrons and therefore different atomic masses A specific type of isotope is also called a nuclide The average atomic mass is the average mass number of all the atoms of a given element including isotopes

Nuclides of Hydrogen Protium (99.9885% of H) = 1p+ Deuterium (0.0115% of H) = 1p+ +1n0 Tritium (negligible in nature, can be synthesized) is radioactive = 1p+ +2n0 How many e- do these isotopes have?

Most isotopes/nuclides don’t have names….. Isotopes are referred to by the element name and the mass number (#p+ + #n0) Ex. Carbon Carbon-14 (14 = 6p+ + 8n0) = radioactive Carbon-12 (12 = 6p+ + 6n0) = most common NOTICE: both types of carbon have same number of protons What if the atom had 7 protons? Practice Problems pg 80

Atomic mass and mass units Metric unit for mass = Too large to measure mass of something so small, so we use amu (atomic mass units) to measure atomic mass Relative atomic mass = the mass of an atom expressed in relation to a defined standard (carbon-12 atom) The carbon-12 atom has a relative mass of 12 amu, so 1 amu = 1/12 mass of C-12 atom

Calculating average atomic mass Depends on mass and relative abundance of isotopes Avg. atomic mass = %A1m1 + %A2m2 + …..

The mole Mole = the amount of a substance that contains as many particles as there are atoms in exactly 12 g of C-12 1 mol of atoms = 6.022 x 1023 atoms Just like… 1 dozen eggs = 12 eggs Avogadro’s number The mole is a counting unit, just like a dozen If all the people on Earth worked to count 1 mole of pennies at 1 penny/second, it would take us all 3 million years to count them!

Molar Mass Mole = the amount (mass) of a substance that contains Avogadro’s number (6.022 x 1023) of particles The mass of 1 mole of a substance = molar mass (found under symbol for each element in PT) I do practice problem 1 on page 85, you do 2-4 alone. We do practice on p 87. handout practice worksheet solo. Assign homework. Be ready for quiz Molar mass for neon = 20.18 g/mol

SUMMARY Atomic # = number of p+ Number of p+ = number of e- Mass number = #p+ + n0 Atomic mass in PT = average atomic mass Units for avg. atomic mass = g/mol 1 mole = 6.022 x 1023 particles e.g. 6.022 x 1023 carbon atoms has a mass of 12.011 g