Chemical Combination Overview, Laws & Examples

Chemical combinations are more than just mad scientist experiments; real science and laws are behind them all. We’ll break down what chemical combination means, the key principles and laws that govern reactions when chemicals are combined, some real-world examples you’ve probably seen before, and even examples of chemical combinations found in nature.

What Is Chemical Combination?

Chemical combination occurs when two or more elements combine to form a new compound with properties different from the original elements. The elements lose their individual identities in the process. The new compound has a fixed composition and properties that are very different from those of the component elements.

Some examples of chemical combinations include:

• Water (H2O) – Hydrogen and oxygen gases combine to form liquid water. The properties of water are very different from hydrogen and oxygen gases.
• Table salt (NaCl) – Sodium metal and chlorine gas combine to form solid table salt. Again, the properties are quite different from the original elements.
• Methane (CH4) – Carbon and hydrogen combine to form methane gas. Methane is a colorless, odorless gas used as a fuel, unlike the black solid carbon and colorless hydrogen gas.

In chemical combination, the elements are bound together by strong chemical bonds to form molecules of the new compound. Ionic, covalent, and metallic bonds are chemical bonds that hold elements together in chemical combinations. The bonds break down the original identities of the individual atoms, and they behave as a whole new substance with different physical and chemical properties.

Factors that Determine a Chemical Combination

The key factors that determine how and whether elements will combine chemically are:

1. Valence electrons – The electrons in the outermost shell that are available for bonding.
2. Electronegativity – Atoms will combine with other atoms that have the appropriate electronegativity difference to form ionic or covalent bonds.
3. Reactivity – Very reactive elements like sodium will readily combine with less reactive elements like chlorine. Less reactive elements are less likely to form compounds.
4. Charge balance – Positively and negatively charged ions will combine to form ionic compounds that have a neutral charge balance.
5. Atomic size- Small atoms have high electron densities and tend to lose or share electrons, while large atoms have low electron densities and tend to gain electrons.
6. Electron configuration – Atoms will gain, lose, or share electrons to achieve a stable outer shell.
7. Molecular geometry – The shape of a molecule based on how atoms or groups of atoms are arranged. Geometry depends on the number of bonds and lone pairs around an atom.

So, in summary, chemical combination results in the formation of new compounds with properties that are distinct from the component elements. The elements are bound together by strong chemical bonds to create a new substance.

Overview of Chemical Combination

Chemical combination, also known as chemical bonding, refers to the joining of atoms together to form molecules. It results in the formation of chemical compounds from elements. The atoms are held together by chemical bonds.

There are three types of strong chemical bonds:

• Ionic bonds: Formed by the electrostatic attraction between oppositely charged ions. For example, sodium chloride (NaCl) consists of Na+ and Cl- ions.
• Covalent bonds: Formed by the sharing of electrons between atoms. For example, hydrogen molecules (H2) and water (H2O) contain covalent bonds.
• Metallic bonds: Formed by the attraction of free-flowing valence electrons to metal ions. For example, iron, copper, and aluminum contain metallic bonds.

Chemical bonds are usually classified as either polar or nonpolar. Polar bonds have an unequal sharing of electrons, resulting in partial charges on the atoms. Nonpolar bonds have an equal sharing of electrons. The polarity of bonds determines whether a molecule is polar or nonpolar overall.

The Key Laws of Chemical Combination Explained

Chemical reactions follow certain rules known as the laws of chemical combination. Understanding these laws can help you better comprehend how elements and compounds interact.

Law of Conservation of Mass

The law of conservation of mass states that for any chemical reaction, the mass of the products equals the mass of the reactants. In other words, mass cannot be created or destroyed in a chemical reaction.

Let’s break this down using a simple example reaction:

Sodium (Na) + Chlorine (Cl2) → Sodium Chloride (NaCl)

• Before the reaction, we have 23 grams of sodium (Na) and 70.9 grams of chlorine (Cl2) for a total mass of 93.9 grams.
• After the reaction, we are left with 58.44 grams of sodium chloride (NaCl).
• The mass before (93.9g) equals the mass after (58.44g) – no mass has been gained or lost.

This holds true for any chemical reaction, as expressed in the following equation:

Mass of Reactants = Mass of Products

Some other key points regarding the law of conservation of mass:

• It applies to both open and closed systems. Whether a reaction happens in a beaker or a test tube, the total mass before and after will be the same.
• It includes all types of matter – solids, liquids, gases, plasma, etc. The form may change, but the total mass remains constant.
• Mass refers to the amount of matter, not the weight. Weight depends on gravity, but mass is an intrinsic property.
• The law only applies to the overall reaction. The mass may change during the reaction process but will return to the original amount once the reaction is complete.
• It is a fundamental law in chemistry and is used to balance chemical equations and calculate reaction yields.

Understanding the law of conservation of mass is key to grasping many other chemical principles. While the types and forms of matter may change during a reaction, the total amount of matter remains the same. This unchanging quality provides a stable foundation for exploring all areas of chemistry.

Law of Definite Proportions

The law of definite proportions, also known as Proust’s law or the law of constant composition, states that chemical compounds are always made up of the same elements in the same proportion by mass. This means that a particular chemical compound always contains exactly the same relative amounts of elements. For example, water (H2O) always contains two hydrogen atoms and one oxygen atom. No matter how much water you have, the ratio of hydrogen atoms to oxygen atoms remains two to one.

Some examples to illustrate this law:

• Ammonia (NH3) always contains nitrogen and hydrogen in a 1:3 ratio. Whether you have a small amount or a large amount of ammonia, it’s always made up of one part nitrogen and three parts hydrogen.
• Sodium chloride (NaCl), or table salt, always contains sodium and chlorine in a 1:1 ratio.
• Carbon dioxide (CO2) always contains one part carbon and two parts oxygen, no matter the quantity.

The law of definite proportions allows scientists to determine the chemical formula of a compound by measuring the masses of its constituent elements. For example, by measuring the masses of hydrogen and oxygen that combine to form water, scientists were able to determine that the formula for water must be H2O.

The law of definite proportions is a fundamental law of chemistry that demonstrates that elements combine in consistent, predictable ways. There are no exceptions to this law. All chemical compounds obey the law of definite proportions, and it applies to all quantities of compounds, both large and small. Understanding this important law paved the way for many discoveries in chemistry and helped scientists develop modern atomic theory.

Law of Multiple Proportions

The law of multiple proportions states that when two elements form a series of compounds, the ratios of the masses of the second element that combine with a fixed mass of the first element will be small whole numbers.

For example, carbon and oxygen form two compounds, carbon monoxide (CO) and carbon dioxide (CO2). In CO, 1 gram of carbon combines with 1.33 grams of oxygen. In CO2, 1 gram of carbon combines with 2.66 grams of oxygen. The ratio of the masses of oxygen that combine with 1 gram of carbon is 1.33:2.66 or 1:2. This illustrates the law of multiple proportions.

The law can be expressed by a mathematical equation. If an element A combines with an element B to form two compounds, the masses of B that combine with a fixed mass of A are in a small whole number ratio.

For compounds: A + bB → AB A + cB → AB2

Then, according to the law of multiple proportions:

b/c = n:1

Where n is a small whole number.

The law of multiple proportions is a result of the fixed characteristic of atoms. Atoms of the same element have the same number of protons and thus the same chemical properties. When elements form compounds, they do so in fixed ratios of small whole numbers in order to satisfy the requirement that compounds must be electrically neutral.

Some examples that illustrate the law of multiple proportions:

• Carbon monoxide (CO) and carbon dioxide (CO2): 1 g C combine with 1.33 g O and 2.66 g O, respectively. The ratio is 1:2.
• Nitrous oxide (N2O) and nitric oxide (NO): 1 g N combines with 1.5 g O and 1 g O, respectively. Ratio is 3:2.
• Methane (CH4) and ethane (C2H6): 1 g C combines with 4 g H and 6 g H, respectively. The ratio is 4:6, which reduces to 2:3.

So, in summary, the law of multiple proportions is a fundamental law of chemistry that establishes the relationship between the masses of elements in a chemical compound. It shows that when two elements form a series of compounds, the ratios of the masses of the second element that combine with a fixed mass of the first element will be small whole numbers.

Gay Lussac’s Law of Gaseous Volumes

Gay-Lussac’s law relates the pressure of a gas to its temperature in a fixed volume. It states that the pressure of a given mass of gas varies directly with the absolute temperature if the volume remains constant. The law can be expressed mathematically as:

P / T = constant or Pi / Ti = Pf / Tf

P/T = k

Where P is the pressure, T is the absolute temperature, and k is a constant. This means that as temperature increases, so does the pressure, and vice versa.

For example, imagine you have a rigid container filled with a fixed amount of gas. As you heat the container, the temperature of the gas increases. The gas molecules gain energy and move faster, colliding with the walls of the container more frequently and with greater force. This results in an increase in pressure. Conversely, cooling the gas causes the molecules to slow down, decreasing the frequency and force of collisions with the container walls, thereby lowering the pressure.

The relationship described by Gay-Lussac’s law holds as long as the amount of gas and the volume remain unchanged. If you change either the amount of gas in the container or the volume of the container, you are changing the conditions, and the pressure-temperature relationship will no longer remain constant.

Gay-Lussac’s law applies to all gases that exhibit ideal gas behavior, meaning the intermolecular forces between molecules are negligible. Most real gases approximate ideal gas behavior under conditions of ordinary pressure and temperature. This law provides a simple model to predict how changes in temperature will affect the pressure of a confined gas. Understanding the relationship between these two variables is fundamental to many applications that involve gases, such as hot air balloons, refrigeration, and internal combustion engines.

In summary, Gay-Lussac’s law states that for a fixed amount of gas at constant volume, the pressure is directly proportional to its absolute temperature. It’s a simple but useful law for predicting the behavior of gases and understanding many practical applications that involve changes in temperature and pressure.

Avogadro’s Law

Avogadro’s law states that equal volumes of gases at the same temperature and pressure contain the same number of molecules. In mathematical terms, it can be expressed as:

V1/n1 = V2/n2

Where V is the volume and n is the number of moles of gas.

This means that the volume of a gas sample is directly proportional to the number of moles of the gas. For example, if you have 4 liters of oxygen gas and 2 liters of hydrogen gas at the same temperature and pressure, the oxygen sample will have twice as many molecules as the hydrogen sample.

Some key points to understand about Avogadro’s law:

• It only applies to gases, not liquids or solids. Gases have free-moving particles, so their volume depends on the number of particles.
• It assumes ideal gases where there are no intermolecular forces between particles. Real gases show some deviation from ideal behavior.
• Volume and number of moles are directly proportional. If volume doubles, moles also double. If volume is cut in half, moles are also cut in half.
• Temperature and pressure must remain constant. Changes in temperature or pressure will affect gas volume, and the relationship will no longer hold.
• One mole of any gas occupies 22.4 liters at standard temperature and pressure (273 K and 1 atm). So you can calculate the number of moles if you know the volume, or vice versa.
• Avogadro’s number (6.022 x 1023 molecules/mole) is an important constant relating moles to actual numbers of molecules or atoms.

By understanding the relationship between volume, moles, molecules, and Avogadro’s number, Avogadro’s law provides crucial insights into the behavior of gases and how their properties can be predicted mathematically. It forms a basis for many applications involving calculations and experiments with gases.

Chemical Combinations in Everyday Life

Chemical reactions happen all around us each and every day. Some of the most common chemical combinations you encounter regularly include:

Baking

When baking, several chemical reactions take place. The combination of wet and dry ingredients creates chemical reactions that leaven and rise, transforming batter into baked goods. Yeast fermentation releases carbon dioxide, causing the bread to rise. The reaction of baking powder or baking soda and an acid like buttermilk or lemon juice also releases CO2. The coagulation of eggs provides structure. All of these chemical reactions work together through baking to create the final product.

Rusting

Rust forms when iron or its alloys, like steel, combine with moisture and oxygen in the air. This chemical reaction creates hydrated iron oxides, commonly known as rust. Rusting occurs in the presence of both water and oxygen and is accelerated by salts like sodium chloride (NaCl). To prevent rusting, you can coat the metal, alloy it with other metals like chromium to form stainless steel, or control exposure to moisture and oxygen.

Photosynthesis

Plants undergo photosynthesis, a chemical reaction where carbon dioxide and water combine in the presence of sunlight and chlorophyll to produce glucose (food) and oxygen. Six molecules of carbon dioxide and six molecules of water combine to form one molecule of glucose and six molecules of oxygen. This vital reaction provides plants with the food they need to grow and thrive and supplies oxygen for animals and humans.

Combustion

Combustion or burning is a fast chemical reaction that produces heat and light. It occurs when a fuel, like wood, coal, or natural gas, combines with oxygen. The products of combustion are carbon dioxide and water vapor. The energy produced by combustion is used for heating, transportation, and generating electricity. By controlling the rate of combustion, we can use it to power engines and generators.

Cellular respiration

Cellular respiration is a chemical reaction in cells that converts nutrients like glucose into energy. Similar to combustion, it combines oxygen and glucose to produce carbon dioxide, water, and energy in the form of ATP. This important metabolic process provides energy for cells to function. In aerobic respiration, oxygen is the final electron acceptor, while in anaerobic respiration, other molecules like lactate or sulfate serve that purpose.

Cooking food

Cooking food often involves chemical reactions like the breakdown of molecules through heating. The application of heat causes proteins to denature, starches to gelatinize and break down, and Maillard reactions to occur. These chemical changes are what give cooked food different flavors and textures than raw ingredients. Boiling, steaming, and grilling are all methods that use heat to chemically transform food.

FAQs About Chemical Combination

Chemical combination, also known as chemical bonding, is a complex topic. Here are some of the most frequently asked questions to help clarify the essential concepts.

What is a chemical bond?

A chemical bond is the force of attraction between atoms in a compound that holds them together. The bond forms when atoms share or transfer electrons to fill their outer electron shells and become more stable. The most common types of bonds are ionic bonds (the transfer of electrons) and covalent bonds (the sharing of electrons).

What are the different types of chemical bonds?

The three main types of chemical bonds are:

• Ionic bonds: The transfer of electrons between metals and nonmetals. Ionic compounds form crystals held together by the attraction between positive and negative ions.
• Covalent bonds: The sharing of electrons between nonmetal atoms. Covalent compounds share electrons to fill their outer shells and become more stable.
• Metallic bonds: The attraction between metal ions and electrons in a metal. The electrons move freely and are shared among the metal ions.

Can chemical bonds be broken?

Yes, chemical bonds can be broken through the absorption of energy. When enough energy is supplied to a compound, its bonds will break, and the atoms or molecules will separate. The amount of energy required depends on the strength and type of bonds in the compound. Stronger bonds, like triple bonds, require more energy to break than weaker single bonds.

What are some examples of chemical combinations?

Some of the most well-known examples include:

• Water (H2O): Formed from two hydrogen molecules and one oxygen molecule.
• Salt (NaCl): Sodium loses an electron to chlorine, forming positively charged sodium ions and negatively charged chloride ions that are attracted to each other.
• Carbon dioxide (CO2): Carbon shares two oxygen atoms, each sharing two electrons.
• Methane (CH4): Carbon shares four hydrogen atoms, each sharing one electron.
• Sucrose (C12H22O11): Carbon, hydrogen, and oxygen combine in a complex molecule to form table sugar.

Conclusion

Understanding chemical combinations is fundamental to chemistry. It helps explain the properties and structure of compounds, as well as many chemical reactions. With practice, you’ll get better at determining the types of bonds present in compounds and predicting molecular geometry.