Chem 103 1st Edition Lecture 6Outline of Last Lecture I. Significant Figuresa. Multiplying and Dividingb. Adding and Subtractingc. Conversiond. Percentage CalculationsII. DensityIII. Atoms and Moleculesa. Symbols and Formulasb. CompoundsOutline of Current Lecture I. The Structure of Atomsa. Protonsb. Neutronsc. ElectronsII. Atomic NumberIII. Mass NumberIV. IsotopesV. Relative massesThese notes represent a detailed interpretation of the professor’s lecture. GradeBuddy is best used as a supplement to your own notes, not as a substitute.VI. Isotopes and Atomic WeightsVII. The MoleCurrent LectureThe Structure of AtomsAtoms are made up of three subatomic particles: protons, neutrons, and electronsAtoms are mostly made up of empty space- To add perspective, if you were a nucleus, your electron would be 9.47 miles awayProtons- Located in the nucleus of an atom- Made of particles called quarks- Carry a +1 electrical charge and have a mass of 1 atomic mass units (u)- The Strong Nuclear Force holds protons to neutrons in nucleusNeutrons- Located in the nucleus- Made up of quarks- Carry no electrical charge and have a mass of 1 atomic mass unit- The Weak Force holds neutrons in the nucleusElectrons- Located outside of the nucleus- Are not made up of any smaller particle- Carry a -1 electrical charge and have a mass of 1/1836 u- Move rapidly around the nucleusAtomic Number of an AtomAtomic number is equal to the number of protons in the nucleus of the atom. In neutral atoms, it’s also the number of electrons. The atomic number is represented by the symbol Z.Mass Number of an AtomMass number is equal to the sum of the number of protons and neutrons in the nucleus. The mass number is represented by the symbol A.IsotopesIsotopes are atoms that have the same number of protons, but a different number of neutrons. They have the same atomic number as their counterparts, but their mass is different. Symbols for isotopes- Represented by the symbol AZE- Z = atomic number, A = mass number- Isotopes can also be represented by notation: Name-A Ex: Magnesium-28- Because Mg has 12 protons, it has 16 neutronsRelative MassesAn atom’s extremely small size makes it inconvenient to use actual masses. Thus relative masses are used. A relative mass is mass that is compared to another mass. Atomic mass units (u) are used to express relative masses of atoms or moleculesAtomic weight is the relative mass of an average atom of the element expressed in atomic mass unitsMolecular weight obtained by adding together the atomic weights of all the atoms in the moleculeEx: The molecular weight of C6H12O6 is 6(12.01) + 12(1.01) + 6 (16.00) = 180.18 uIsotopes and Atomic WeightsIn nature, elements often occur as a mixture of isotopes. Because of this, the element’s atomic weight is the average mass of the mixture of isotopes. This is generally done by using percentages of isotopes found in the mixture of isotopes. Calculating atomic weight:-- This equation can be used to find the atomic weight of a mixture of isotopes. The result of the equation should agree with that found on the periodic tableThe MoleAtomic mass numberMg2812Atomic numberOne mole (also called Avogadro’s Number) = 6.022 x 1023When referring to one mole of an element, it is worth 6.022 x 1023 atoms of the element. One mole sample of one element is the same number of atoms as a one mole sample of another element. However, one mole an element can weigh differently than one mole of a different element.It should be kept in mind that grams and amu’s are different:- One mole Na = 22.09 g Na- One atom Na = 22.09 u NaThe mole concept can apply to molecules and compounds as well- Ex: 1 mol C3H8O3 molecules3 mol C 8 mol H 3 mol O- How many H atoms are in 2.5 mol of H2SO4?2.5 mol H2SO4 x (2 mol H atoms) x (6.022 x 10 23 H atoms) = 3.011 x 1024 atoms H(1 mol H atoms)(1 mol