CHEMISTRY

1. Differentiate between elements, compounds and molecules

 

Elements, compounds, and molecules are fundamental concepts in chemistry. Here's how they differ:

1. Elements:
• An element is a pure substance made up of only one type of atom.
• Each element is represented by a unique chemical symbol, such as "H" for hydrogen, "O" for oxygen, and "Fe" for iron.
• Elements cannot be broken down into simpler substances by chemical means. They consist of identical atoms with the same number of protons in their nuclei.
2. Compounds:
• A compound is a substance composed of two or more different elements chemically combined in fixed ratios.
• Compounds have properties different from those of their constituent elements. For example, water (H2O) is a compound made up of hydrogen and oxygen, and it has different properties than pure hydrogen or oxygen.
• Compounds can be broken down into their constituent elements through chemical reactions.
3. Molecules:
• A molecule is the smallest unit of a compound or element that retains the chemical properties of that substance.
• Molecules can be made up of either one type of atom (as in O2, molecular oxygen) or multiple types of atoms (as in H2O, water).
• Molecules can exist as discrete entities and are formed through covalent bonding, where atoms share electrons.

 

 

2. Describe the structure of the atom in terms of protons, neutrons and electrons.

 

Structure of an Atom: An atom is the basic unit of matter, and its structure consists of three main subatomic particles: protons, neutrons, and electrons.

1. Protons:
• Protons are positively charged subatomic particles found in the nucleus of an atom.
• They have a relative mass of approximately 1 atomic mass unit (amu) and carry a charge of +1.
2. Neutrons:
• Neutrons are neutral subatomic particles also located in the nucleus.
• They have a relative mass of approximately 1 amu and carry no electric charge (neutral).
3. Electrons:
• Electrons are negatively charged subatomic particles that orbit the nucleus in specific energy levels or electron shells.
• They have a much smaller mass, about 1/1836 amu, and carry a charge of -1.
• Electrons are involved in chemical reactions, as they are responsible for forming bonds between atoms to create molecules.

The number of protons in the nucleus defines the element, and the sum of protons and neutrons in the nucleus gives the atom's atomic mass. Electrons determine the chemical behavior of an atom, as they are involved in interactions with other atoms to form compounds and molecules. The arrangement of electrons in the electron shells or energy levels is essential in understanding an atom's chemical properties and reactivity.

 

 

3. Determine the number of subatomic particles in an atom form an element's atomic number and mass

We can determine the number of subatomic particles in an atom, specifically the number of protons, neutrons, and electrons, using the element's atomic number and atomic mass (also known as mass number). Here's how to do it:

1. Protons (Z - Atomic Number): The atomic number (Z) of an element represents the number of protons in the nucleus of each atom of that element. For example, if the atomic number is 6, as in the case of carbon (C), there are 6 protons in each carbon atom.
2. Neutrons (Mass Number - Atomic Number): The number of neutrons in an atom can be calculated by subtracting the atomic number (protons) from the mass number (atomic mass). The mass number is typically given as a whole number on the periodic table or in the atomic symbol of an element. For example, if the atomic mass of carbon is 12, and the atomic number is 6, then the number of neutrons is 12 - 6 = 6.
3. Electrons (Same as Protons for a Neutral Atom): In a neutral atom, the number of electrons is equal to the number of protons. This maintains a neutral charge for the atom. So, if there are 6 protons in a carbon atom (as determined by the atomic number), there are also 6 electrons.

Keep in mind that the number of protons (atomic number) uniquely identifies the element, and variations in the number of neutrons result in different isotopes of the same element. Isotopes have the same number of protons (same element) but different numbers of neutrons, resulting in variations in atomic mass.

 

4. Write the electronic configuration of an atom using Bohr model

The Bohr model of the atom is a simplified representation that describes the electron arrangement in energy levels or electron shells. In this model, electrons orbit the nucleus in distinct energy levels, with each energy level accommodating a specific maximum number of electrons. The Bohr model is useful for understanding basic electron arrangement, but it has limitations and has been replaced by more accurate quantum mechanical models. Here's how you can write the electronic configuration of an atom using the Bohr model:

1. Identify the Atomic Number: Determine the atomic number of the element you want to represent. The atomic number (Z) indicates the number of protons in the nucleus, which is equal to the number of electrons in a neutral atom.
2. Assign Electrons to Energy Levels: In the Bohr model, electrons are distributed into energy levels, which are labeled with numbers (n) starting from the innermost level (n = 1) and increasing outward. Each energy level can hold a specific maximum number of electrons: 2n².
3. Fill the Energy Levels: Place electrons into energy levels, starting from the lowest energy level (n = 1) and moving to higher levels. Follow the "2n²" rule to determine the maximum number of electrons each level can hold. Distribute electrons according to this rule.
4. Observe the Octet Rule (for Some Elements): For the outermost energy level (valence shell), try to arrange electrons to achieve a stable octet (eight electrons). This rule is particularly applicable to the elements in the first two rows of the periodic table, as they tend to follow the octet rule.

For example, let's write the electronic configuration of carbon (C) using the Bohr model:

• Carbon has an atomic number of 6, which means it has 6 electrons.
• In the Bohr model, we start filling the energy levels. The first energy level (n = 1) can hold a maximum of 2 electrons, and the second energy level (n = 2) can hold a maximum of 8 electrons.
• Place 2 electrons in the first energy level and 4 electrons in the second energy level.

So, the Bohr model electronic configuration for carbon (C) is:

1st energy level (K-shell): 2 electrons 2nd energy level (L-shell): 4 electrons

This configuration represents the simplified electron arrangement in a carbon atom using the Bohr model. Keep in mind that this model doesn't capture the full complexity of electron behavior, which is better explained by quantum mechanics.

 

 

5. Identify metals and non-metals on the periodic table.

In the periodic table, elements are categorized into different groups based on their properties. Two primary categories of elements are metals and non-metals. Here's how to identify metals and non-metals on the periodic table:

Metals:

1. Location: Metals are found on the left and in the middle of the periodic table.
2. Physical Properties: Metals typically exhibit the following physical properties:
• Shiny luster (when freshly polished)
• Good conductors of heat and electricity
• Malleable (can be hammered into thin sheets)
• Ductile (can be drawn into wires)
• Solid at room temperature (with the exception of mercury, which is a liquid)
3. Chemical Properties: Metals tend to lose electrons in chemical reactions, forming positively charged ions (cations). They generally react with non-metals to form ionic compounds.

Common examples of metals include iron (Fe), copper (Cu), gold (Au), and aluminum (Al).

Non-Metals:

1. Location: Non-metals are typically found on the upper right side of the periodic table, including elements in Groups 14, 15, 16, and 17.
2. Physical Properties: Non-metals exhibit the following physical properties:
• Lack of metallic luster (often appear dull)
• Poor conductors of heat and electricity
• Brittle (not malleable or ductile)
• Varied states at room temperature (can be solid, liquid, or gas)
3. Chemical Properties: Non-metals tend to gain or share electrons in chemical reactions, forming negatively charged ions (anions) or covalent compounds.

Common examples of non-metals include hydrogen (H), oxygen (O), nitrogen (N), and carbon (C).

It's important to note that there is also a category known as metalloids, which are elements with properties that fall between those of metals and non-metals. Metalloids are typically found in a diagonal "staircase" region on the periodic table, separating the metals from the non-metals. Common metalloids include silicon (Si), germanium (Ge), and arsenic (As).

 

 

6. Predict ionic charges from the periodic table

Predicting ionic charges from the periodic table involves understanding the concept of valence electrons and the octet rule. Valence electrons are the electrons in the outermost energy level (valence shell) of an atom. The number of valence electrons influences the charge an atom is likely to gain or lose to achieve a stable electron configuration, often following the octet rule (having 8 electrons in the outermost shell).

Here are some general guidelines for predicting ionic charges based on an element's position in the periodic table:

1. Alkali Metals (Group 1): Elements in Group 1, such as sodium (Na) and potassium (K), have one valence electron. They tend to lose this electron to achieve a stable configuration, resulting in a charge of +1.
2. Alkaline Earth Metals (Group 2): Elements in Group 2, like magnesium (Mg) and calcium (Ca), have two valence electrons. They tend to lose these two electrons to form ions with a charge of +2.
3. Halogens (Group 17): Elements in Group 17, including fluorine (F) and chlorine (Cl), have seven valence electrons. They tend to gain one electron to complete their octet, resulting in a charge of -1.
4. Noble Gases (Group 18): Noble gases, such as helium (He) and neon (Ne), have a full complement of valence electrons and are stable. They typically do not form ions and have a charge of 0.
5. Transition Metals: Transition metals, located in the middle of the periodic table, often have multiple possible ionic charges. The charge of a transition metal ion depends on the specific compound and its chemical context. Roman numerals are used to indicate the charge of the transition metal in compounds. For example, Fe²⁺ and Fe³⁺ represent iron ions with charges of +2 and +3, respectively.
6. Non-Metals: Non-metals, which are primarily on the right side of the periodic table, tend to gain electrons to achieve a full valence shell. The charge they acquire depends on the number of electrons they gain. For example, oxygen (O) gains two electrons to form O²⁻ with a charge of -2.
7. Metalloids: Metalloids, like silicon (Si) and germanium (Ge), can exhibit a range of charges, depending on the specific compound they form. They often follow similar charge patterns to non-metals.

These are general trends, and there can be exceptions based on specific chemical reactions and contexts. When working with ionic compounds, it's important to understand the charges of the ions involved and balance them to achieve a neutral overall charge in the compound.

 

 

7. Define ions and differentiate cations from anions

Ions are electrically charged particles formed when atoms gain or lose electrons. Atoms are electrically neutral because the number of positively charged protons in the nucleus is balanced by the number of negatively charged electrons orbiting the nucleus. However, when an atom gains or loses electrons, it acquires a net electrical charge, becoming an ion.

Cations and anions are two types of ions with different charges:

1. Cations:
• A cation is a positively charged ion.
• Cations are formed when an atom loses one or more electrons.
• This loss of electrons results in an excess of protons, giving the ion a net positive charge.
• Common cations are typically formed by metals, which tend to lose electrons to achieve a stable electron configuration. For example, a sodium atom (Na) loses one electron to become a sodium cation (Na⁺), which has a +1 charge.
2. Anions:
• An anion is a negatively charged ion.
• Anions are formed when an atom gains one or more electrons.
• This gain of electrons results in an excess of electrons, giving the ion a net negative charge.
• Common anions are typically formed by non-metals, which tend to gain electrons to achieve a stable electron configuration. For example, a chlorine atom (Cl) gains one electron to become a chloride anion (Cl⁻), which has a -1 charge.

In summary, cations are positively charged ions formed by the loss of electrons, typically associated with metals, while anions are negatively charged ions formed by the gain of electrons, typically associated with non-metals. The charges of these ions are crucial in understanding ionic compounds, where cations and anions come together to form neutral compounds by balancing their charges.

 

8. How to differentiate between neutral atoms, isotopes and ions

Differentiating between neutral atoms, isotopes, and ions involves understanding their fundamental characteristics and how they differ in terms of composition and electrical charge. Here's how to distinguish between these three concepts:

1. Neutral Atoms:
• Composition: Neutral atoms consist of a nucleus in the center, composed of protons and neutrons, surrounded by electrons in electron shells or energy levels.
• Electrical Charge: A neutral atom has an equal number of protons (positively charged) and electrons (negatively charged). The positive and negative charges balance, resulting in an overall charge of 0.
• Example: A neutral hydrogen atom has one proton, one electron, and no charge (0).
2. Isotopes:
• Composition: Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons. They have the same atomic number but different atomic masses.
• Electrical Charge: Isotopes are still neutral; they have the same number of protons and electrons, which cancel out to give an overall charge of 0.
• Example: Carbon has two stable isotopes, carbon-12 (12C) and carbon-13 (13C). Both are neutral atoms with the same number of protons and electrons, but they differ in the number of neutrons.
3. Ions:
• Composition: Ions are atoms (or groups of atoms) with an unequal number of protons and electrons. They can be positively charged (cations) or negatively charged (anions).
• Electrical Charge: Cations have more protons than electrons, resulting in a positive charge, while anions have more electrons than protons, leading to a negative charge.
• Example: A sodium ion (Na⁺) is a cation with 11 protons and 10 electrons, giving it a net positive charge of +1. A chloride ion (Cl⁻) is an anion with 17 protons and 18 electrons, resulting in a net negative charge of -1.

In summary, neutral atoms have an equal number of protons and electrons, isotopes are variants of an element with different numbers of neutrons while still being neutral, and ions have an unequal number of protons and electrons, resulting in a net electrical charge. Understanding these distinctions is fundamental to comprehending the behavior and properties of atoms and molecules in chemistry.

 

 

9. Draw electron transfer diagrams for simple ionic compounds

 

10. Write chemical formulas for simple ionic compounds

 

11. Explain the Law of Conservation of Mass using experimental data

 

12. Identity and explain the difference between a physical change and a chemical change

 

13. Apply the Law of Conservation of Matter to balance chemical equations.