Chem lecture Ch. 2 PDF

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Lecture notes on chapter 2 general chem ...

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General Chemistry 1 Chapter 2 Part 1 Atomic Structure ●

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Nucleus ● Protons (+ charge) ● Neutrons (no charge) Electrons (- charge)

Subatomic Particles ● ●

Protons and neutrons have very similar masses (~ 1 amu) Electrons have a much lower mass (~ 0.0005 amu)

Atomic Number ● ● ●

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The atomic number, denoted by Z, is equal to the number of protons in the nucleus of an atom Determines the identity of an atom (see periodic table for atomic number & symbol locations) Examples: ● Mg (magnesium): Z = ? ● Sn (tin): Z = ? In an electrically neutral atom, the number of electrons is equal to the atomic number (# protons = # electrons)

Mass Number ● ● ● ●

Ions

The mass number, denoted by A, is equal to the sum of the protons and neutrons in an atom Mathematically: A – Z = n (where n is the number of neutrons) Nuclide notation is a shorthand method for designating an element and its mass number and atomic number Element symbol is written, with the mass number A as a superscript to the left of the symbol and the atomic number Z as a subscript (also to the left of the symbol)

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Atoms can gain or lose electrons to form ions that are not electrically neutral (ions carry positive or negative charge) Ions do not have the same number of electrons as protons (number of electrons ≠ Z)

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Types of ions Cation: # protons > # electrons (+ charge)

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Anion: # electrons > # protons (- charge

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Examples of neutral atoms and corresponding ions: S, S2Cation or anion? Difference between # protons and # electrons? Al, Al3+ Cation or anion? Difference between # protons and # electrons?

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Isotopes are forms of the same element that differ in the number of neutrons within the nucleus They still have the same number of protons (same atomic number) Many elements have more than one isotope Ex: C-12, C-13, and C-14 Some of the isotopes are energetically unstable and can give off energy and particles to achieve stability (radioactivity)

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Isotopes

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Atomic Mass Units (“amu”) ● ● ● ●

The standard used to determine the masses of the atoms in the periodic table Based on the most abundant isotope of carbon: 12C = exactly 12 atomic mass units (amu) , 1 amu = 1/12 the mass of one carbon-12 atom Relationship between amu and grams: 1 amu = 1.661 x 10-24 g

Calculating Atomic weight ●

The average atomic weight for an element is a weighted average of its isotopes

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Mathematically: Isotope-1 mass (% abundance/100) + Isotope-2 mass (% abundance/100) + …

Format of the Periodic Table ● ● ● ● ● ● ●

Arranged in groups and periods Columns are the groups (elements in the same group have similar physical & chemical properties) Main-group elements Transition metals Inner transition metals Rows are the periods The heavy stair-step line on the right side distinguishes between metals, nonmetals and metalloids

Metals ● ● ● ● ● ●

To the left of the heavy stair-step line on the periodic table Common characteristics: lustrous (shiny) malleable (can be shaped or bent) ductile (can be stretched into wire) conduct heat and electricity

Nonmetals ● ● ● ● ●

To the right of the heavy stair-step line on the periodic table Common characteristics: not lustrous break easily insulate against heat and electricity

Metalloids ● ●

These are the elements that form the stair-step boundary between metals and nonmetals (B, Si, Ge, As, Sb, Te) They have some characteristics of metals and some of nonmetals

Molecular Compounds ● ●

Molecules are distinct substances made up of 2+ atoms linked together by covalent bonds (sharing of electrons between nuclei) Examples:

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● Sulfur trioxide, SO3 ● Dinitrogen tetroxide, N2O4 The molecular formula indicates the number of each type of atom present in one molecule of a compound

Empirical Formula ●

An empirical formula is the ratio in which elements are found to be present in compounds (regardless of the molecular structure). The subscripts are the smallest possible whole-number ratios.

Ionic Compounds ● ● ●

Often formed from the combination of metals and nonmetals Involves transfer of electrons from one atom to another The ions must combine in a ratio that makes the compound electrically neutral overall

Polyatomic Ions ● ● ●

Polyatomic ions contain two or more atoms Can carry an overall positive or negative charge Behave like monoatomic ions

Common Polyatomic Ions ● ●

Cation: NH4+ Anions: SO42-, MnO4-, C2O42-, PO43-

Naming Binary Ionic Compounds ● ● ●

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Write the name of the first element Write the stem name of the second element, changing the ending to the suffix –ide If the metal component is one that can have more than one type of charge (ex: Cu can be +1 or +2), put the appropriate charge in Roman numerals in parentheses behind the metal’s name No prefixes are needed (prefixes are only used in naming binary covalent compounds)

The Mole ● ●

The mole is the basic counting unit used to indicate how many atoms, molecules, formula units or ions are in a given substance. 1 mole corresponds to 6.022x1023 particles; this is “Avogadro’s Number” (NA)

The Mole: Elements ●

For any element, we can write:

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● 1 mol = atomic mass in g ● 1 mol = 6.022x1023 atoms ● Example ● 1 mol of P = 30.97g P ● 1 mol of P = 6.022x1023 P atoms ● These relationships can be used in dimensional analysis to interconvert between g, mol, and number of atoms

Molar Mass ●

The average molar mass of a molecule is the sum of the atomic masses of the elements in that molecule.

Formula Mass ● ● ●

Formula mass is a more general term than molar mass. It can be applied to molecular compounds or ionic compounds. It is the sum of the atomic masses of the elements in a molecule or an ionic compound. The term “molar mass” technically only applies to molecules, since ionic compounds do not exist in the form of individual molecules.

The Mole: Compounds ●

For any molecular (covalent) compound, we can write: ● 1 mol = molar mass in g ● 1 mol = 6.022x1023 molecules ● For any ionic compound, we can write: ● 1 mol = formula mass in g ● 1 mol = 6.022x1023 formula units

Mass Percent ●

The mass percent of an element in a compound can be calculated using the following equation: Mass percent = total mass component/ total mass whole substance x 100%

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This is also sometimes referred to as the “percent by mass”.

Empirical & Molecular Formulas

The empirical formula for a compound is the simplest whole number ratio of the atoms in the compound. As we have already seen, the molecular formula tells us the actual number of each element in a molecule of the compound. The molecular formula might be the same as the empirical formula, OR it might be a whole number multiple of the empirical formula.

Determining Molecular Formulas ●

Given the molecular weight for a compound and its empirical formula, we can determine its molecular formula.

Hydrated Compounds ●

Hydrates are compounds in which molecules of water are associated with the ions of the compound

Determining Hydrate Formulas ●

A hydrate sample with a known mass is heated to drive off all of the water. The sample is weighed after heating to determine how much anhydrous salt remains and how much water was lost. From the masses of anhydrous salt and water, the formula for the hydrate can be determined: ● Convert both masses to moles ● Divide the moles of water by the moles of anhydrous salt; this mole ratio gives you the “x” in the general hydrate formula (salt ∙ xH2O)...