Chapter 3

Elements and Compounds

All known substances on Earth and probably the universe are formed by combinations of more than 100 elements. Most substances are compounds, composed of two or more elements. Compounds can be decomposed into two or more simpler substances. Water can be decomposed into hydrogen and oxygen.

Table salt can be decomposed into sodium and chlorine.

An element is a fundamental or elementary substance that cannot be broken down into simpler substances by chemical means. Each element has a number.

Beginning with hydrogen as 1, the elements are numbered in order of increasing complexity.

An atom is the smallest particle of an element that can exist. It is the smallest unit of an element that can enter into a chemical reaction. A molecule is a combination of two or more atoms. O2 is an element. If a molecule contains only atoms of one kind, it is an H2O is a compound. element. Otherwise, it is a compound. Symbols of the elements: One or two letters. First letter always capitalized. Elements are not distributed equally by nature.

Carbon Nitrogen Barium Sodium

C N Ba Na

In the universe, the most abundant element is hydrogen (91%) followed by helium (8.75%). All other atoms combined are present in 0.25%.

Oxygen is the most abundant element in the crust of the earth (49.2%), followed by silicon. Oxygen is the most abundant element in the human body (65%), followed by carbon. But carbon is < 0.1% in earth crust!

A number of symbols appear to have no connection with the element. Most symbols start with the same letter as the element.

Hydrogen is in a group of its own

leftmost transition area rightmost

metals

nonmetals

Metals, Nonmetals and Metalloids Highly reactive metals form minerals by Metals Most elements are metals. combining with other elements in nature. Physical properties Good conductors of heat and electricity. Nonmetals 17 elements are nonmetals, including Have high luster. rare and radioactive (At) and six noble High melting point and density. gases. Remaining ten nonmetals make Solid at room temperature. almost all compounds. Mercury is the only exception.

Malleable (hammered into sheets); Ductile (drawn into wires). Mix with other metals to form alloys Brass = copper + zinc Bronze = copper + tin Steel = Iron + carbon Chemical properties Little tendency to combine and form compounds with other metals Readily form compounds with nonmetals (especially oxygen). Some less reactive metals are found in nature in free state (elemental form).

Physical properties: opposite of metals Lack luster. Poor conductors.

Low melting point. Low density.

Can be found in all three phases (solid, liquid, gas) at room temperature. C, P, S, Se, I: solids Br: liquid H, O, N, F, Cl, noble gases: gases Metalloids

only seven elements

Be, Si, Ge, As, Sb, Te, and radioactive Po. Physical properties:

Intermediate

Periodic Table The periodic table was designed by Dimitri Mendelev in 1869 In the table each element’s symbol is placed inside of a box Above the symbol of the element is its atomic number. Elements are arranged in order of increasing atomic number Elements with similar chemical properties are organized in columns called families or groups.

7

N He

These elements are known as the noble gases. They are non-reactive. They are found in pure (elemental) form.

Ne

Noble metals can also be found in elemental form. Most substances around us are mixtures or compounds.

Ar

Air is a mixture of nitrogen, oxygen, argon and traces of other gases (e.g. water vapor and carbon dioxide).

Kr

H, O, N etc. cannot exist in nature as single atoms. They are always paired (H2, O2, N2). These elements exist as diatomic molecules.

Xe Rn

Composition of Elements and Compounds Hydrogen and oxygen are elements. Elements cannot be decomposed into simpler substances. Free, or elemental hydrogen (H2) is found in volcanoes; It can be prepared in the laboratory. Elemental oxygen (O2) is found in air (21 vol%); It can be prepared in the laboratory.

Water is a compound; its formula is H2O. It is made of two hydrogen atoms and one oxygen atom. Water contains no free hydrogen nor free oxygen.

H2O

A compound is a distinct substance that contains two or more elements combined in a definite proportion by weight.

H2

O2

Compounds can be decomposed chemically into simpler substances – that is, into simpler compounds or elements.

Atoms of the elements that constitute a compound are always present in simple whole number ratios. They are never present as fractional parts. Physical and chemical properties of a compound differ from those of the parent elements.

Types of Compounds

Ionic compounds are made of ions.

Compounds can be molecular or ionic. Molecular compounds are made of atoms.

An ion is a positively or negatively charged atom or group of atoms. Cation is positively charged ion.

A molecule is the smallest uncharged individual unit of a compound formed by the union of two or more atoms. Chemical bonds that keep atoms joined in molecules are called covalent bonds.

Anion is negatively charged ion.

Ionic compounds are held together by electrostatic attractive forces between positively and negatively charged ions. Table salt, sodium chloride is a colorless crystalline ionic substance, 39.3% sodium and 60.7% chlorine by mass. The solid NaCl does not conduct electricity; molten NaCl, and solution of salt in H2O do.

Compounds can be classified as molecular or ionic. Molecular compounds are held together by covalent bonds. Ionic compounds are held together by attractive forces between their positive and negative charges.

Chemical Formulas

Serve as abbreviations of the names of compounds. Tell which elements the compound is composed of and how many atoms of each element are present in a formula unit.

Show the symbols of the atoms of the elements present in a compound.

CaCl (1) 2

Show the ratio of the atoms of the elements present in a compound.

calcium chloride

calcium chlorine

Ca calcium Cl chlorine

2 Cl 1 Ca

Rules for Writing Chemical Formulas Formulas do not necessarily represent the arrangement of atoms. When a formula contains one atom of an element, the symbol of that element represents the one atom. The number one (1) is not used as a subscript. When the formula contains more than one atom of an element, the number of atoms is indicated by a subscript written to the right of the symbol of that atom. When the formula contains more than one of a group of atoms that occurs as a unit, parentheses are placed around the group, and the number of units of the group is indicated by a subscript placed to the right of the parentheses. There is one phosphorus atom in a phosphate group

Ba3(PO4)2 There are three barium atoms

Indicates two phosphate (PO4)3- groups

There are four oxygen atoms in a phosphate group

There are three barium, two phosphorus and eight oxygen atoms in the compound!

H2O

NaCl Indicates the element sodium (one atom)

Indicates the element chlorine (one atom)

H3PO4 Indicates the element hydrogen Indicates 3 H atoms

Indicates the element oxygen Indicates 4 O atoms

Indicates the element phosphorus (1 atom)

Chapter 4

Properties of Matter

Properties of a substance are characteristic of the substance. Each substance has a set of properties that are characteristic of that substance and give it a unique identity. Physical properties are inherent characteristics of a substance. They can be determined without changing its composition.

Color Taste / Odor Physical state Melting point Boiling point

Chemical properties describe the ability of a substance to form new substances, either by reaction with other substances or by decomposition. Physical properties of chlorine gas (Cl2): 2.4 x heavier than air yellowish-green color disagreeable odor melting point: -101 oC boiling point: -34.6 oC

Chemical properties of chlorine (Cl2): Doesn’t burn supports combustion Bleaching agent disinfectant Produces table salt with sodium

Physical and Chemical Changes Physical changes are changes in physical properties (such as size, shape, and density) or changes in the state of matter without an accompanying change in composition. No new substances are formed.

tearing of paper change of ice into water change of water into steam heating platinum wire

Chemical change is a change in which new substances are formed that have different properties and composition from the original material. The formation of copper(II) oxide from copper and oxygen is a chemical change. The newly-formed substance has set of properties that are different from those of copper. Water decomposes when electrical energy passes through it. The products, elemental hydrogen and oxygen, have different properties from those of water.

Chemical Reaction Water decomposes into hydrogen and oxygen when electrolyzed. Chemical symbols can be used to express chemical reactions

reactant

2H2O

yields

products

2H2 + O2

Conservation of Mass No change is observed in the total mass of the substances involved in a chemical change.

sodium + 46.0 g

sulfur

→

32.1 g

78.1 g reactant mass reactants

sodium sulfide 78.1 g

→

78.1 g product

=

mass products

Energy

Kinetic energy matter possesses due to its motion.

Mechanical Chemical Electrical Heat Nuclear Radiant

Potential energy could also thought as the energy stored in an object. Chemical energy is one form of potential energy. The heat released when gasoline burns is associated with a decrease in its chemical potential energy. The new substances formed by burning have less chemical potential energy than the gasoline and oxygen.

Increasing potential energy

Potential energy is the energy that an object possesses due to its relative position.

Bouncing ball. Running man.

Height

Energy is the capacity to do work.

Potential energy rises with the height increase

Energy in Chemical Reactions In all chemical changes, matter either absorbs or releases energy. Examples of chemical processes leading to energy release Type of Energy

Energy Source

Electrical

Storage batteries

Light

A lightstick. Fuel combustion.

Heat and Light

Combustion of fuels.

Body

Chemical changes within body cells.

Examples of chemical processes leading to energy absorption Type of Energy

Chemical Change

Electrical

Electroplating of metals. Decomposition of water into hydrogen and oxygen

Light

Photosynthesis in green plants.

Conservation of Energy An energy transformation occurs whenever a chemical change occurs.

If energy is given off in a chemical change, the products will have less chemical potential energy than the reactants.

If energy is absorbed during a chemical change, the products will have more chemical potential energy than the reactants. Energy can be neither created nor destroyed, Law of Conservation of Energy though it can be transformed from one form of energy to another form of energy.

higher potential energy

Electrolysis of Water

lower potential energy

Burning of hydrogen in air

Heat: Quantitative Measurement Heat is a form of energy associated with small particles of matter. Temperature is a measure of the intensity of heat, or of how hot or cold a system is. The SI unit for heat energy is the joule (pronounced “jool”). Another unit is the calorie.

1 calorie = 4.184 Joules (exactly)

This amount of heat energy will raise the temperature of 1 gram of water by 1 oC. Two calories is needed to raise 1 gram of water by 2 oC. Twice as much heat energy is required to raise the temperature of 200 g of water by 10 oC as compared to 100 g of water. specific Heat = heat

(

q = C

)(

x

mass

m

100 Ag water

200 Bg water

30oC 20

30oC 20

4184 J

8368 J

temperature change

)( x

)

∆t

Specific Heat The specific heat of a substance is the quantity of heat required to change the temperature of 1 g of that substance by 1 oC. The units of specific heat in calories are:

 calories   gram oCelcius   

 cal   g oC   

The units of specific heat in joules are:

Joules    gram oCelcius   

 J   g oC   

It takes about 5 times more energy to warm 1 g of water by 1 oC as to warm the same amount of aluminum by 1oC.

Quantitative Calculation of Heat The relation of mass, specific I can be found if heat, temperature change both sides were (Δt), and quantity of heat lost divided by m x ∆t or gained is expressed by the general equation:

(

)(

specific heat of substance

Problem 1: Calculate the specific heat of a solid in J/goC and if 1638 J raise the temperature of 125 g of the solid from 25.0oC to 52.6oC.

)

mass of Δt = heat substance heat = 1638 J mass = 125 g ∆t = 52.6 oC – 25.0 oC = 27.6 oC spec.heat = ?

heat specific heat =

mass x ∆t

=

1638 J 125 g x 27.6 oC

= 0.475

J g oC

Heat = specific heat Q

C

=

x

mass of water m

x

x x

change in temp. ∆t

Problem 2: How much heat (in J) it takes to warm a 2,000. lb block of iron from 50 oC to 75 oC? J 453.6 g m = 2,000. lb x 25 oC = 1.07 x 107 J Q = m C ∆t = 2,000. lb x x 0.473 g oC 1 lb ∆t = 25 oC = 1.1 x 107 J (2 significant figures!)

From Table 4.3 (p.71): C = 0.473 J/goC

Problem 3: What is the final temperature of a 20.0 g block of iron after it has absorbed 100.0 J of heat, if its initial temperature was 25.0 oC? Hint: solve for ∆t, then add it to the initial temp! ∆t =

Q mC

100.0 J =

J 20.0 g x 0.473 g oC

Calorie Measurement Place the food (e.g. Twinkie) into a vessel (“bomb”), surround it with water and measure the water temperature as you burn the food.

= 10.6 oC


Final temp. = 25.0 oC + 10.6 oC = 35.6 oC

Advanced A sample of a metal with a mass of 212 g is heated to 125.0 oC and then problem: dropped into 375 g of water at 24.0 oC. If the final temperature of the water is 34.2 oC, what is the specific heat of the metal? When the metal enters the water, it begins to cool, losing heat to the water. At the same time, the temperature of the water rises. This process continues until the temperature of the metal and the temperature of the water are equal (34.2 oC), at which point no net flow of heat occurs.

Heat gained by water = heat lost by metal mW = 375 g

∆tW = 34.2 – 24.0 oC

Strategy: ∆tM = 125.0 – 34.2 oC 1. Calculate the heat (qW) gained by the water. mM = 212 g 2. Calculate the final temperature of the metal. spec.heat o 3. Calculate the specific heat of the metal. of water = 4.184 J/g C qW = qM temperature gained by water 1.

∆tW = 34.2

oC

– 24.0

oC

= 10.2

3. oC

heat gained by water J x 10.2 oC = 1.60 x 104 J qW = 375 g x 4.184 g oC

qM = qW = 1.60 x 104 J qM spec.heat = m x ∆t 1.60 x 104 J spec.heat = 125 g x 90.8 oC

temperature lost by metal 2.

∆tM = 125.0

oC

– 34.2

oC

= 90.8

oC

spec.heat = 0.831

J g oC

Homework, chapter 3 (paired exerc., p.59): 5. Write the formula for each compound: a) Zinc oxide (1 atom Zn, 1 atom O) b) Potassium chlorate (1 atom K, 1 atom Cl, 3 atoms O) c) Sodium hydroxide (1 atom Na, 1 atom O, 1 atom H) d) Ethyl alcohol (2 atoms C, 6 atoms H, 1 atom O) 11. How any total atoms are represented in each formula? a) Co(ClO3)2; b) (NH4)2SO3; c) CH3CH2COOH; d) C12H22O11. 19. Classify each material as an element, compound or mixture, as per the picture representing three boxes filled with various gases (right). 23. What percent of the first 36 elements on the periodic table are metals? Homework, chapter 4 (paired exerc., p.76): 7. State whether each of the following represents a chemical or physical change: a) A steak is cooked until

well done; b) Students fire-polish the end of a glass rod to smoothen jagged edges; c) Chlorine bleach removes coffee stains from white lab coat; d) When two clear solutions are mixed together, the solution becomes yellow and cloudy. 13. Indicate with the plus sign (+) any of these processes that requires energy, and a negative sign (-) any that release energy: a) melting ice; b) starting a car; c) flash of lighting; d) dry ice changing to vapor; e) blowing up a baloon. 19. A 135 g metal bar requires 2.50 kJ to change its temperature form 19.5 oC to 100.0 oC. Calculate the specific heat of the metal. 20. A 275 g of metal requires 10.75 kJ thus raising its temperature from 21.2 oC to its melting point (327.5 oC). Calculate the specific heat of antimony. 30. A sample of copper was heated to 275.1 oC and placed into 272 g of water at 21.0 oC. The temperature of the water rose to 29.7 oC. How many grams of copper were in the sample (specific heat of copper is 0.385 J/goC)?