Intro chemistry Chpt 5 Chemical Accounting

Percent by volume:

If both solute/sovent are liquids " " is often used because liquid volumes are easily measured. % by volume= volume of solute/volume of solution x 100%

Molarity (M):

Is the amount of solute , in moles per liter of solution (M)= moles of solute/liters of solution. Solution concentration: molarity/moles Calculate the molarity of a solution made by dissolving 3.50 mol NaCl in enough water to produce 2.00 L of solution Solution: Molarity (M)=moles of solute/liters of solution= = 3.50 mole NaCl/2.00 L solution = 1.75 M NaCl We read 1.75 M NaCl as "1.75 molar Sodium chloride"

Aqueous solutions:

Those in which water is the solvent

Chemical equation:

Uses symbols/formulas to represent the elements and compounds involved in the change The plus sign (+) indicates that carbon and oxygen are added together or combined in some way. The arrow (→) is read "yield(s)" or "react(s) to produce." Substances on the left of the arrow (C and O₂ in this case) are reactants, or starting materials Those on the right (here, CO₂) are the products of the reaction. Reactants/products need not be written in any particular order in a chemical equation, except that all reactants must be to the left of the arrow and all products to the right. Example: carbon reacts with oxygen to form carbon dioxide C + O₂ → CO₂ In other words, we could also write the preceding equation as O₂ + C → CO₂ At the atomic or molecular level, the equation means that one carbon (C) atom reacts with one oxygen (O₂) molecule to produce one carbon dioxide (CO₂) molecule.

Dilute solution:

One that contains a relatively small amount of solute in lots of solvent. For example: a pinch of sugar in a cup of tea makes a dilute, faintly sweet, sugar solution

Concentrated solution:

One with relatively large amount of solute dissolved in a small amount of solvent. For example: Pancake syrup, with lots of sugar solute in a relatively small amount of water, is a concentrated, and very sweet, solution

Stoichiometric factor:

Relates the amounts, in moles, of any 2 substances involved in a chemical reaction. Consider the combustion of butane, a common fuel for lighters 2 C₄H₁₀ +13 O₂→ 8 CO₂ + 10 H₂0 The coefficients in the balanced equation allow us to make statements such as: * 2 mol C₄H₁₀ react with 13 mol O₂ * 8 mol CO₂ are produced for every 2 mol C₄H₁₀ that react * 10 mol H₂O are produced for every 8 mol CO₂ produced We can turn these statements into conversion factors know as " " ***Note*** that we can set up " " for any 2 substances involved in a reaction. Also note that " " because they come from the balanced equation for a reaction, involve small whole numbers.

Avogadro's Number:

6.02 x 10²³ Avogadro had no way of knowing how many molecules were in a given volume of gas. Scientists since his time have determined the number of atoms in various weighed samples of substances. Atoms are so incredibly small that these numbers are extremely large, even for tiny samples. Recall that the mass of a carbon-12 atom is exactly 12 atomic mass units (u) because the carbon-12 atom sets the standard for atomic mass. The number of carbon-12 atoms in a 12-g sample of carbon-12 is called Avogadro's number And has been determined experimentally to be 6.0221367 x 10²³. We usually round this number to three significant figures: 6.02 x 10²³

The Mole:

" A dozen eggs and a mole of sugar please " We buy socks by the pair (2 socks), eggs by the dozen (12 eggs), pencils by the gross (144 pencils), and paper by the ream (500 sheets of paper). We use these quantitative terms as constant measures/are comfortable with them. But a dozen eggs/a dozen oranges Do Not weigh the same. If an orange weighs 5 times as much as an egg, a dozen oranges will weigh 5 times as much as a dozen eggs. Furthermore, when we need to obtain large numbers of these items, we can do so easily No one would want to purchase 2000 individual sheets of paper it it meant counting them out, but buying 4 reams of paper is easily done. Likewise, a baker might need to make hundreds of cakes/therefore need many dozens of eggs/can obtain them in multiples of a dozen. Sellers of many of those items will not want to sell them in fragments, however, so even if staff in an office wanted to buy 500 pencils, they would not typically be able to buy the exact number. They would instead have to purchase 4 gross of pencils (576 pencils) and have a surplus. Similarly, sometimes we have to deal with an excess amount of chemical components due to the whole-number ratios used in chemical equations. The same way that an office manager may purchase paper by the ream/pencils by the gross, a chemist counts atoms/molecules by the mole. A single carbon atom is much too small to see or weigh, but a mole of carbon atoms fills a tablespoon and weighs 12 g. A mole of carbon and a mole of titanium each contain the same number of atoms. But a titanium atom has a mass 4 times that of a carbon atom So a mole of titanium has a mass 4 times that of a mole of carbon.

Although simple equations can be balanced by trial/error, a couple of strategies often help:

1). If an element occurs in just one substance on each side of the equation, try balancing that element first. 2) Balance any reactants or products that exist as the free element last 3). When you add a coefficient to a compound to correct one element, be aware that it will also change the number of atoms in other related elements The most important step in balancing any equation is to check that the results you obtain is indeed balanced. Remember that for each element, the same number of atoms of the element must appear on each side of the equation.

2 common pitfalls in the process of balancing equations, as well as the correct method:

1). You cannot simply gain or lose substances on the product side of the equation, as it may represent an invalid reaction. 2). Also, if you change the subscript of a specific chemical formula, you create either a different substance or and invalid formula.

Percent composition of a compound from formula masses:

A chemical formula represents the ratio of atoms as well as a ratio of mass. This relationship allows us to be able to relate the mass to the number of particles/be able to express the information in a meaningful way for a variety of applications. For example: The mass of 1.00 mol CO₂ is 44.0 g, of which 12.0g is carbon And 32.0 g is oxygen The composition of CO₂ can be expressed as percentages by mass of carbon/of oxygen: % by mass C= mass C/mass CO₂ x 100% = 12.0 g/44.0 g x 100% = 27.3% Because oxygen is the only other element in CO₂, its percent by mass is: 100% - 27.3% = 72.7% oxygen

Solution:

A homogeneous mixture of 2 or more substances.

How are atoms and molecules the same and how are they different?

All the elements are listed out in a periodic table. If it's in the table, it's an element! Atoms can join together - they form bonds together - to make MOLECULES. For example, two atoms of hydrogen hook together to form a molecule of hydrogen, H₂ for short. Note 2 self, I had to put this info in because it gets confusing. It will help me clear up the confusion 4 myself hopefully.

A mole (mol):

An amount of any substance or item that contains the same number of elementary units as there are atoms in exactly 12 g of carbon-12. That number is 6.02 x 10²³ Avogadro's number The elementary units may be atoms (such as S or Ca), molecules (such as O₂ or CO₂), ions (such as K⁺ or SO₄²⁻) or any other kind of formula unit. A mole of NaCl, for example: contains 6.02 x 10²³ NaCl formula units, which means that it contains 6.02 x 10²³ Na⁺ ions and 6.02 x 10²³ Cl⁻ ions

Avogadro's hypothesis:

Based on shrewd interpretation of experimental facts, was that equal volumes of all gases, when measured at the same temperature/pressure, contain the same number of molecules This means that if we weigh equal volumes of different gases, the gases won't weigh the same. The ratio of their masses should be the same as the mass ratio of the molecules themselves The challenge is that there is not always the option of using gases only when performing reactions. We need to extend this relationship to other forms of matter.

Formula Mass:

For any substance, the " " is the sum of the masses of the atoms represented in the formula If the formula represents a molecule, the term molecular mass is often used. For example, because the formula CO₂ Specifies one carbon atom/2 oxygen atoms per molecule of carbon dioxide, the formula (or molecular)mass of carbon dioxide (CO₂) is the atomic mass of carbon plus twice the atomic mass of oxygen. Formula mass CO₂: =1 x atomic mass of C+2 x atomic mass of O =1 x 12.0 u + 2 x 16.0 u = 44.0 u

Balancing chemical equations

Many chemical reactions require more thought for example, hydrogen reacts with oxygen to form water. Using the formulas we can represent this reaction as. However, this representation shows 2 oxygen atoms in the reactants (as O₂), but only one in the product in (H₂O). H₂+O₂→ H₂O (not balanced) Because matter is neither created nor destroyed in a chemical reaction (the law of conversation of mass) The equation must be balanced to represent the chemical reaction correctly. Because we are interpreting the equation to mean that each symbol represents multiples of a single atom or molecule, we cannot simply assume that we are using one-half of the oxygen molecule. Instead , we will need to maintain the relationship between the whole particles. Also, we must be sure that the same number of each of the different types of atom appears on both sides of the arrow. To balance the oxygen atoms, we need only place the coefficient 2 in front of the formula for water. This coefficient means that 2 whole molecules of water are produced. As is the case with a subscript of 1, a coefficient of 1 is understood when there is no number in front of the formula. Everything in a formula is multiplied by the coefficient in front of it. In the preceding equation, adding the coefficient 2 before H₂O not only increases the total number of oxygen atoms on the product side to 2 but also increases the total number of hydrogen atoms to 4. But the equation is still not balanced. There are now 4 hydrogen atoms on the product side of the equation but only 2 on the left side of the equation. As we fixed the number of oxygen atoms in the product, we unbalanced the number of hydrogen atoms. H₂+O₂→2H₂O (not balanced) To balance the hydrogen, we must now place a coefficient 2 in front of H₂ Now there are 4 hydrogen atoms/2 oxygen atoms on each side of the equation. Atoms are conserved: The equation is balanced and the law of conservation of mass is obeyed. 2H₂+O₂→2H₂O (balanced)

Percent by mass:

Many commercial solutions are labeled with the concentration in percent by mass. For example, sulfuric acid is sold in several concentrations: 35.7% H₂SO₄ for use in storage batteries, 77.7% H₂SO₄ for the manufacture of phosphate fertilizers and 93.2% H₂SO₄ for pickling steel. Each of these percentages is by mass: 35.7 g of H₂SO₄ per g of sulfuric acid solution and so on. % by mass= mass of solute/mass of solution x 100% As with percent by volume, we can express the masses of solute/solution in any mass unit, as long as we use the same unit for both.

Molar mass:

The " " of a substance is just what the name implies: The mass of 1 mole of that substance. The molar mass is numerically equal to the atomic mass, molecular mass, or formula mass, but it is expressed in the: unit grams per mole (g/mol). The atomic mass of sodium is 23.0 u So its molar mass is 23.0g/mol The molecular mass of : carbon dioxide is 44.0 u So its molar mass is 44.0 g/mol The formula mass of: ammonium sulfate is 132.1 u So its molar mass is 132.1 g/mol We can use these facts, together with the definition of the mole, To write the following relationships: 1 mol Na = 23.0 g Na 1 mol CO₂ = 44.0 g CO₂ 1 mol (NH₄)₂SO₄ = 132.1 g (NH₄)₂SO₄ These relationships supply the conversion factors we need to convert between mass in grams and amount in moles, as illustrated in the fooling examples. It is important to not that these conversion factors represent equalities Which means that we can always invert them in the equations so that the units cancel appropriately.

Law of combining volumes:

The law states that when all measurements are made at the same temperature and pressure, the volumes of gaseous reactants/products are in small whole-number ratios.

Stoichiometry:

The quantitative relationship between reactants and products in a chemical reaction. The ratio of moles of reactants and products is given by the coefficients in a balanced chemical equation Consider the combustion of butane, a common fuel for lighters 2 C₄H₁₀ + 13 O₂→ 8 CO₂ + 10 H₂O

Solute:

The substance being dissolved.

Solvent:

The substance doing the dissolving. The solvent is usually present in greater quantity that the solute. Examples of solvents: Kerosene dissolves grease ethanol dissolves many drugs isopentyl acetate, a component of banana oil, is a solvent for model airplane glue water is the most familiar solvent, dissolving as it does many common substances such as sugar, salt and ehtanol