Chem 1151: Ch. 2 Atoms and Molecules

Structure of the Atom Mass (g)Mass (u) Proton (p+)1.67 x Neutron (n)1.67 x Electron (e-)9.07 x /1836 Most of the mass is actually in the nucleus Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

The origin of atoms Following the big bang  Expansion of space  Cooling  Formation of fundamental particles  Formation of low mass nuclei (H and He)  Star formation  Fusion reactions forming heavier elements

Synthesis of Matter Nucleosynthesis: Protons and neutrons join to form nuclei. Fusion: Multiple nuclei join to form heavier nucleus. Formation of heavier elements: Two protons collide – Releases positron, neutrino Nucleus with proton and neutron collides with another proton – releases gamma ray Two He-3 atoms collide – Produces He-4 – Releases two protons

Periodic Table of the Elements All matter in our universe categorized in periodic table. – Based on atomic number (number of protons). Arranged in columns (groups) and rows (periods). – Groups have similar properties. – Periods correspond to filling of quantum shells by electrons.

Elements from Group 7A chlorine bromine iodine Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

Elements You Should Know

Applications of Atomic and Mass Numbers – On the periodic table, the atomic number is written as a whole number above the symbol F. – In the written description, fluorine is said to have 9 protons (the atomic number is the number of protons). – In the symbol, the number 9 is written in the atomic number or Z (lower left) position. – Note: The periodic table does not show the mass number for an individual atom. It lists an average mass number for a collection of atoms! Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

Isotopes Isotopes are atoms that have the same number of protons in the nucleus but different numbers of neutrons. That is, they have the same atomic number but different mass numbers. Because they have the same number of protons in the nucleus, all isotopes of the same element have the same number of electrons outside the nucleus. Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

Isotope Symbols Isotopes are represented by the symbol, where Z is the atomic number, A is the mass number, and E is the elemental symbol. Isotopes are also represented by the notation: Name-A, where Name is the name of the element and A is the mass number of the isotope. An example of this isotope notation is magnesium-26. This represents an isotope of magnesium that has a mass number of 26. Since all of the mass of atom comes from the protons and the neutrons, the mass number is the total number of protons and neutrons. You can therefore determine the number of neutrons by subtracting the atomic number from the mass number.

Use of Elemental Notation Q: How to represent lead-208? Q: How many p+, e-, n?82, 82, 126 Q: How to represent element X with 4 p+ and 5 n.

Relative Masses and Mass Units The extremely small size of atoms and molecules makes it inconvenient to use their actual masses for measurements or calculations. Relative masses are used instead. Relative masses are comparisons of actual masses to each other. For example, if an object had twice the mass of another object, their relative masses would be 2 to 1. An atomic mass unit is a unit used to express the relative masses of atoms. One atomic mass unit is equal to 1/12 the mass of a carbon-12 atom. A carbon-12 atom has a relative mass of 12 u because carbon- 12 has 6 protons and 6 neutrons. Mass (g)Mass (u) Proton (p+)1.67 x Neutron (n)1.67 x Electron (e-)9.07 x /1836 Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

Determining Mass However, carbon has an actual mass listed of u, not 12 u. Why are these values different? C Mass (g)Mass (u) Proton (p+)1.67 x Neutron (n)1.67 x Electron (e-)9.07 x / Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

How Isotopes Determine Atomic Weight The atomic weight of an element is the relative mass of an average atom of the element expressed in atomic mass units. Many elements have more than 1 isotope (e.g. – 12 C, 13 C, 14 C). Abundance of isotopes are not evenly distributed. Weighted atomic mass of Carbon ( 12 C, 13 C only) = ( *12u) + ( * u) = u. C C AMU12 u u

Determining Atomic Weight A specific example of the use of the equation is shown below for the element boron that consists of 19.78% boron-10 with a mass of u and 80.22% boron-11 with a mass of 11.01u. This calculated value is seen to agree with the value given in the periodic table. Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

Molecular Weight The relative mass of a molecule in atomic mass units is called the molecular weight of the molecule. Because molecules are made up of atoms, the molecular weight of a molecule is obtained by adding together the atomic weights of all the atoms in the molecule. The formula for a molecule of water is H 2 O. This means one molecule of water contains two atoms of hydrogen, H, and one atom of oxygen, O. The molecular weight of water is then the sum of two atomic weights of H and one atomic weight of O: MW = 2(at. wt. H) + 1(at. wt. O) MW = 2(1.01 u) + 1(16.00 u) = u Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

Molecular Weight The clear liquid is carbon disulfide, CS 2. It is composed of carbon (left) and sulfur (right). What is the molecular weight for carbon disulfide? Answer: MW = 1(atomic weight C) + 2(atomic weight S) u + 2(32.07 u) = u Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

The Mole THE MOLE CONCEPT APPLIED TO ELEMENTS – The number of atoms in one mole of any element is called Avogadro's number and is equal to 6.022x – A one-mole sample of any element will contain the same number of atoms as a one-mole sample of any other element. – One mole of any element is a sample of the element with a mass in grams that is equal to the atomic weight of the element. EXAMPLES OF THE MOLE CONCEPT – 1 mole Na = g Na = 6.022x10 23 Na atoms – 1 mole Ca = g Ca = 6.022x10 23 Ca atoms – 1 mole S = g S = 6.022x10 23 S atoms Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

The Mole THE MOLE CONCEPT APPLIED TO COMPOUNDS – The number of molecules in one mole of any compound is called Avogadro's number and is numerically equal to 6.022x – A one-mole sample of any compound will contain the same number of molecules as a one-mole sample of any other compound. – One mole of any compound is a sample of the compound with a mass in grams equal to the molecular weight of the compound. EXAMPLES OF THE MOLE CONCEPT – 1 mole H 2 O = g H 2 O = 6.022x10 23 H 2 O molecules – 1 mole CO 2 = g CO 2 = 6.022x10 23 CO 2 molecules – 1 mole NH 3 = g NH 3 = 6.022x10 23 NH 3 molecules Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

Relationships: Mass, Moles, Molecular Weight 1 mol C atoms = x atoms C x atoms C = g C 1 mol C atoms = g C C has atomic weight of u or g/mol 1 mol S atoms = x atoms S x atoms S = 32.1 g S 1 mol S atoms = 32.1 g S S has atomic weight of 32.1 u or 32.1 g/mol

Problems 1. What is the mass in g of 1.35 mol of S? Mass S? Mol  g 1.35 mol S X32.1 g S =43.3 g S 1 mol S 2. How many S atoms are in 98.6 g of S? 98.6 g S X 1 mol S X x atoms S = 1.85 x atoms S 32.1 g S1 mol S 3. What is the mass in g of 1 atom of S? 1 atom SX 1 mol S X 32.1 g = 5.33 x g S x atoms S 1 mol S

Moles of Molecules 1. What is the mass in g of 1.62 mol of O 2 molecules? MW of O 2 = 2 x (atomic weight of O) = 2 x (16.0 u) = 32.0 u 1 mol O 2 molecules = x molecules O x molecules O 2 = 32.0 g O 2 1 mol O 2 molecules = 32.0 g O mol O 2 molecules X 32.0 g O 2 = 51.8 g O 2 1 mol O 2 molecules 1 mol O 2 molecules = x atoms O

Compound (Molecular) Formulas The compound (molecular) chemical formula represents the numerical relationships that exist between atoms in a compound. This also applies to moles. Compound formula: all elements and number of each in a compound Examples: Urea1C, 4H, 2N, 1O Hydrofluoric acid 1H, 1F Sodium bicarbonate2H, 1C, 3O Sodium Azide 1Na, 3N 1 molecule of H 2 SO 4 contains 2 atoms of H 1 atom of S 4 atoms of O 1 mol of H 2 SO 4 contains 2 mol of H 1 mol of S 4 mol of O

Mole Calculations (continued) The mole concept applied earlier to molecules can be applied to the individual atoms that are contained in the molecules. An example of this for the compound CO 2 is: 1 mole CO 2 molecules = 1 mole C atoms + 2 moles O atoms g CO 2 = g C g O 6.022x10 23 CO 2 molecules = 6.022x10 23 C atoms + (2) 6.022x10 23 O atoms Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

Finding the Molecular Weight *Use the mole relationship of the compound to find the MW Ex. 1 H 2 SO 4 Atomic Weights H = g/mol S = g/mol O = g/mol MW of 1 mol of H 2 SO 4 = (2 * H = g/mol) + (1 * S = g/mol) + (4 * O = g/mol) = g/mol Ex. 1 C 3 H 8 O (isopropyl alcohol) Atomic Weights C = g/mol H = g/mol O = g/mol MW of 1 mol of C 3 H 8 O = (3 * C = g/mol) + (8 * H = g/mol) + (1 * O = g/mol) = g/mol

Finding the Number of Atoms in a Compound Ex. 1 How many C atoms in 5.2 g of C 3 H 8 O (isopropyl alcohol) Atomic Weights C = g/mol H = g/mol O = g/mol MW of 1 mol of C 3 H 8 O = (3 * C = g/mol) + (8 * H = g/mol) + (1 * O = g/mol) = g/mol 1 mol of C 3 H 8 O contains 3 mol of C

Percent Composition Mass relationships can be used to determine percent compositions 1 mol of H 2 SO 4 % H = 2.0 gx100=2.0 % 98.1 g % S = 32.1 gx100=32.7 % 98.1 g % O = 64.0 gx100=65.2 % 98.1 g

% Mass of element in compound % mass N in HNO 3 ?N = g/mol H = g/mol O = g/mol HN0 3 % mass= part X 100 total Total= (1 x H) + (1 x N) + (3 + O) = (1 x 1.008) + (1 x 14.01) + (3 x 16.00) = g/mol HN0 3 % mass= g/mol X 100 = 22.23% g/mol

% Mass of element in compound % mass N in NaN 3 ?N = g/mol Na = g/mol NaN 3 % mass= part X 100 total Total= (1 x Na) + (3 x N) = (1 x g/mol) + (3 x g/mol) = g/mol NaN 3 % mass= g/mol X 100 = 64.64% g/mol

Atomic Weight of Element with Multiple Isotopes Element X has 3 isotopes(10)X 70% u (11)X 20% u (12)X 10% u What is atomic weight of element X? (70% x m) + (20% x u) + (10% x u) = (70 x m) + (20 x u) + (10 x u) = u 100

Ions  Ion: Atom or molecule that has either lost or gained electrons from valence shell resulting in a net charge (positive or negative) compared to the number of protons.  Ions of element have the same number of protons, but a different number of electrons.  For example compare the following:  Common atomic ions you should know:  H +, Na +, K +, Mg 2+, Ca 2+, Fe 2+, Fe 3+, Ag 1+, Pb 2+,  N 3-, P 3-, O 2-, S 2-, F -, Cl -, Br - vs.

Exercises What are the charge, mass (u) in the following nuclei? 1.5 p+, 6 n 2.10 p+, 12 n 3.11 p+, 14 n, 1e- What are the p+, n, e-, charge? What are the MW values in u? C 6 H 12 O 6 (glucose) O 3 (ozone) Which is denser, U-235 or U-238? Why?

What these numbers mean? Mass of 1 atom of Mg = x g Mg atomic weight = u How many atom of Mg in24.31 g Mg? g Mg x 1 atom Mg = x atoms Mg x g Mass of 1 atom of C = x g C atomic weight = u g C x 1 atom C = x atoms C x g The number of atoms of an element, in a mass equivalent to it’s atomic weight, is equal to Avogadro’s number (mole). or 1 u = 1 g/mole